Chemistry Unit I
Beginning of the Universe: “big-bang” theory - before the beginning of the universe, only quarks existed
due to the extreme temperature, density, and pressure. Then, universe experienced an explosion of
subatomic particles and energy 14 / 13.7 billion years ago, and is expanding ever since.
- Evidence: 1) Doppler Effect & Hubble’s Law - Frequency increases with direction of travel. Ex.
train whistle pitches.
- Red shift: electromagnetic radiation emitted by stars in distant
galaxy would appear to be shifted downward in frequency
- Blue shift: stars moving closer appear to be shifted upward in
frequency
- Hubble’s Law: size of the red shift is proportional to the distance and speed of the star
moving away
- 2) Nucleosynthesis: creation of atomic nuclei when the universe’s temperature was at 1 billion
degrees, deuterium, lithium, and helium
- 3) Cosmic Microwave Background: Microwaves which look like those radiated by a black body >0K
degrees
Nebula: debris left over from star formation that can accrete into planets, moons, asteroids
Planet Formation: gas & dust attracted by gravity and electrostatic forces. Heated by radioactive decay large portion melts. Ni and Fe separated because of differential crystallization.
k
Summaries: Due to natural forces, nebulas form interstellar bodies, heat/pressure/etc. Are necessary for
uneven distribution of elements and are responsible for Earth’s current composition.
Early Atomic Structure: Greeks (400 B.C.) continuous theory of matter (matter can keep dividing) vs.
discontinued theory of matter (a smallest unit exists)
1) Aristotle and Plato - continuous
2) Democritus - existence of smallest individual, “atomos”, individual in Greek
John Dalton: (1776 - 1844) “Father of Chemical Atomic Theory” - first to public comprehensive atomic
theory.
1) Dalton’s theory published over the years (1803 - 1807)
a) Chemical elements made up of tiny, indivisible, indestructible units called atoms: these
particles maintain identity through chemical and physical changes only rearranged
A Lavoisier: Law of Conservation of Mass ( matter is not created nor destroyed), “Father of Quantitative
Chemistry”
Element: A chemical substance resisting simplification by chemical means
- The atoms of an element are identical (modified because of isotopes)
- Atoms of different elements - different
Basic Atomic Theory cont.: atoms combine in fixed, small whole-number ratios
- Coincides w/ Joseph Proust’s law of definite proportions
- States that compounds are composed of elements combined in a fixed ratio by mass
Other work by Dalton:
1) Law of Multiple Proportions: the ratio of a fixed mass of one element when combined with the
multiple masses of another element may be expressed in small whole number
a) No fraction of atoms
Dalton’s Model: Spherical, indivisible, indestructible (1807)
a) We can not directly observe atoms, however we can image atoms on the surface of
materials utilizing instrumentation such as the STM
i)
Resolution on the order of a nanometer or less (10^-9)
Nanotechnology: Molecular science- has applications with STM, can be used to move atoms individually
as well as to generate high res maps of material surfaces
Summaries:
Electricity:
1) Static Electricity: build-up of electrical charge on an object (recognize most objects are electrically
neutral)
a) Being static implies that it is not moving
2) These objects of one charge are attracted to objects of the opposite charge, objects with the same
charge repel each other
a) Electrostatic force: force between charged objects
3) Electrically charged objects will return to their neutral condition. Why?
a) Enthalpy: stability increases when energy decreases, stability decreases, when energy
increases
b) Entropy: state of disorder, stability increases when entropy increases, stability decreases
when entropy decreases
4) This flow of electrical charge is referred to as Electrical current
a) Electrical current measured in units of amperes- electrical current which transports one
coulomb of charge past a point in one second.
5) The behavior of electrical charges is often observed as electrodes
a) Electrode- site or terminal of electrical charge
i)
2 types of electrodes:
(1) Anode
(2) Cathode
Coulomb’s Law:
Proportionality constant- allows variables to function
together
Force inversely proportional to square of the distance
between charges
Force is proportional to the product of the magnitudes
of the charges
Force is attractive for charges of opposite signs and
repulsive for charges of the same sign
Particles and Fields:
1) a charged particle moving through a field may have its path altered by the field
a) It will be attracted toward the opposite pole, and repelled by the like pole
b) Particle deflected if the field is sufficiently weak, however strong fields may result in the
capture of the particle. Deflection by a magnetic field at a right angle.
Alpha & Beta Particles:
a) Alpha particles: helium nuclei- no electrons, positive charge
b) High speed electrons, negative charge
Voltage:
1) The strength of a field or current is related to the potential energy of the electrons or quantity of
the charges present. The potential difference is measured in units of volts- the push behind
electrons movement.
a) The volt is defined as a potential difference of one joule/coulomb
b) The greater the voltage, the greater the potential energy of the system, and the more work
that may be done by that system
c) Majority of household volts: 120 - 240 if electric stove/dryer
2) Experiment frequently requires evacuated or partially evacuated containers (metal/glass) as
gaseous molecule present may interfere with particles.
3) Studied through interactions: fluor/photographic film
CRT (Cathode Ray Tube):
1) Geissler 1855 develops mercury pumps and decent vacuum tubes
2) Hittorf improved vacuum, first to see cathode ray
3) Plucker showed cathode ray altered by a magnet
4) Hittorf placed solid body in front of cathode, cuts off glow from walls of tube, moves in straight
line
5) Varley suggest made of particles
6) Crookes incorrectly assumes them to be molecules of gas with a negative charge from the cathode
7) Goldstein ray originates at cathode, names light emitted as cathode ray
a) Observes the CRT producing radiation that moves toward cathode- canal rays
Observations using the CRT:
1) The cathode ray can push a small paddle wheel up an incline against gravity
a) Ray carried energy and could do work
2) The ray is deflected from the straight line path using a magnetic field
a) Possible relationship not yet discovered
Thomson discovers the electron:
1) Deflects cathode ray with an electric field
a) Rays bend towards positive poles
i)
Ray must have a negative charge (1897 corpuscles [electrons] are discovered to be
negatively charged)
2) Also determined the charge to mass ratio of the electron
8
a) 𝑒/𝑒 = 1.759 × 10
𝑒/𝑒
b) Millikan performed the oil drop experiment in 1909 to determine charge of the electron
i)
Results showed that the value of 𝑒𝑒 varied but was always a multiple of a small
number
(1) Ionized- to be electrically charged
ii)
Resulted from variation in ionizations via x-rays in lower chambers
iii)
Millikan determined the terminal velocity of the droplet in order to calculate its
mass, some electrons strike the droplets, giving them a charge
iv)
6𝑒𝑒𝑒𝑒 is the viscous resistance due to air described by Stoke’s law for a spherical
drop of radius (r) and moving through air with a velocity of v and a viscosity of𝜂
v)
In the absence of a field the drop equals the force due to the resistance of air
−19
(1) 1.602 × 10
𝑒for plates with strength to E to opposed the 𝑒𝑒 gravitational
force, the drop of charge remains stationary
(2) Believed to be the smallest possible charge
3) The Plum Pudding model:
a) J.J. Thomson developed an atomic model (including electrons) called plum pudding model,
modifications:
i)
Atom is divisible
ii)
Atom has an electrical character/nature
iii) Canal ray: the rays eventually found to be composed of
the ionized gas molecules in the cathode ray tube
S
Summaries:
Proton discovery:
1) In 1886, Eugene Goldstein observed luminous rays streaming through a cathode whose
modification was that it had holes- fritted cathode
a) It also contained hydrogen gas at low pressure
2) The magnitude and direction of the rays in the presence of a magnetic field indicated they were
positive, had a mass significantly larger than the electron, and travel at a much lower velocity
3) Credit for the discovery of the proton is debatable pending upon the source
a) However, some credit Rutherford with its discovery
4) The proton’s mass was eventually identified as 1.67265 × 10
−24
𝑒
X rays:
1) Wilhelm C. Röntgen (1845 - 1923)
a) ≈1895- accidentally discovered x-rays when his wife passed her hand in the path of x-rays,
revealing her bone structure and wedding ring. This discovery inspired Becquerel to study
fluorescence and phosphorescence
i)
Worked with potassium uranyl sulfate
b) Radioactivity- the spontaneous emission of subatomic particles of energy by disintegration
of atomic nuclei
Types of Radiation:
1) Ionizing potential, penetration, and composition
a) Alpha Radiation:
i)
ii)
iii)
iv)
particles carry charge of +2 and have a mass of ~6.65 × 10
alpha somewhat poor penetrating radiation
symbol 4 2 𝑒𝑒
strong ionizing potential
−24
𝑒
b) Beta Radiation:
i)
composed of high speed electron particles carry charge of -1 and have a mass of
9.11 × 10−28 𝑒
ii)
beta has fair penetrating ability
iii)
equation symbol is 0 −1 𝑒sky
iv)
good ionizing potential
c) X-Rays:
i)
they have good penetrating ability
ii)
poor ionizing potential
(1) more energy, increased ionizing potential
d) Gamma Rays:
i)
Similar to x rays (electromagnetic radiation) but usually carry more energy
(1) They have very strong penetrating ability
(2) Symbol is y
(3) Variable ionizing potential
Gold Foil Experiment
1) Rutherford (1871 - 1973) along with coworkers Geiger, Marsden, and Bohr devised and performed
the gold foil experiment in 1911
a) The experiment: a thin sheet of gold foil was enclosed by a fluor covered circular shroud
b) There was an opening in the shroud through which alpha particles were shot at the gold
foil
2) Results:
a) The vast majority of alpha particles passed through the foil without any deflection from
their path
b) Some of the particles passed through the foil with a minor deflection to their path
c) A scant few (1 in 8000) alpha particles were deflected at large angles, in essence backwards
d) Most alpha particles passed through the foil without deflection because the vast majority of
the atom is empty space with a few electrons in it
i)
The few alpha particles deflected at minor angles came close to a small bundle or
core of positive charge which were, therefore repelled
ii)
The scant number of alpha particles being bounced back was a result of an almost
direct collision with a very massive (extremely dense) and positively charged region
occupying a very small molecule
Atomic Models:
1) Dalton’s atomic model
a) Spherical indivisible unit
2) Thomson- the plum pudding model
a) A sea of positive charge with electrons scattered throughout
b) A divisible atom - subatomic particles existed
c) An electrical nature associated with the atoms
3) Rutherford - planetary model
a) The atom contained a small, positively charged, dense core or center called the nucleus
b) The electrons traveled around outside the nucleus
c) Most of the atoms volume was actually empty space
d) The nuclear diameter is about 10−4 the diameter of the atom
Problems with the Planetary Model:
1) Why weren’t the electrons pulled into the positively charged nucleus of the atom
a) Responded by stating that the electrons motion prevents it from being pulled into the
nucleus much the same as the planets aren’t pulled into the sun
b) According to classical mechanics charged particles moving in a curved path should emit
energy or some other form of electromagnetic radiation. Eventually, they would lose
enough energy to be pulled into the nucleus
Neutron:
1) In 1928, a german physicist, Walter Bothe, and his student, Herbert Becker took the initial step in
the search for the neutron. They bombarded beryllium with alpha particles emitted from
polonium and found that it gave off a penetrating, electrically neutral radiation, which they
interpreted to be high-energy photons. Eventually, the neutron was discovered in 1932 when
James Chadwick used scattering data to calculating the mass of this neutral particle.
a) Chadwick is credited with discovery
Basic Structure of Atomic particles:
a) The proton with a standardized charge of +1, a mass 1.673 x 10−24 g, and located in the
nucleus
b) The neutron has a charge of 0, a mass of 1.678 𝑒 10−24 g, and located in the nucleus
i)
The neutron is slightly larger than the proton
c) The electron with a standardized charge of -1. A mass of 9.11 x 10−28g, and located outside
the nucleus in what is called the electron cloud
i)
The electron is ~1/1837 the mass of the proton
Facts About the Basic Structure of the Modern Atom:
1) Atoms are usually found in their neutral state, having equal numbers of protons and electrons
a) # protons = # electrons for NEUTRAL atoms
2) Most of the atom is composed of the empty space (vacuum), the volume of the atom being
established by the electron cloud
3) The nucleus is an extremely small, extremely dense, positive core of the atom which contains the
vast majority of the
4) atoms mass
a) Nuclear density is about 1013 to 1014 g/𝑒𝑒3
5) Why don’t the electrons follow Newton’s first law and fly off and out of the atom
a) Electrostatic force- coulombic attraction
6) Atomic diameters are expressed on the order of angstroms
a) 1 angstrom = 10−10 m
b) symbol for angstrom = Å
7) Atomic nuclei have diameters that are on the order of 10−4 angstroms
8) The identity of the element is established by the number of protons contained in the nuclei of its
atoms the identity of an element is established by the number of protons contained in the nuclei of
its atoms
a) This number of protons is referred to as the atomic number of the element
b) The symbol for the atomic number = Z
9) Atoms may lose or gain electrons without any change in element identity, BUT any change in
number of protons will create a new identity for that atom
a) Loss of electrons = cations
b) Gain of electrons = anions
Isotopes:
1) Atoms of the same element which differ in mass this mass difference is due to possession of
differing number of neutrons
a) Recognize, nuclear charge cannot change
2) Isotopes of an element have a different mass numbers
a) Mass numbers (symbol A) is equal to the sum of major nucleons
b) A nucleon refers to any particle in the nucleus of an atom
Mass Number:
1) Therefore, mass numbers equals the sum of protons and neutrons
a) Mass number = Z + #n or A = Z +#n
b) Since A is the symbol for mass # to find the number of neutrons in an isotope use A - Z = #n
Representing Isotopes:
1) Isotopes have the general
2) X is the chemical symbol of
3) If the isotope is an ion, a
form of
the element
charge will be in the right superscript
Naming Isotopes:
1) The term nuclide is used to refer to a specific atom of an isotope
a) An atom with specific number of protons and neutrons
2) Hydrogen has 3 isotopes with special names:
a) Hydrogen-1: protium, Hydrogen-2: deuterium, Hydrogen-3: tritium
3) Other naming involves the element name, a dash, and the mass number
Atomic Mass Scale:
1) Atoms are too small to measure their mass directly, however their relative mass may be
determined compounds are evaluated for the combining ratios of the atoms after which upon
decomposition, an analysis of the mass of each element may be established
2) From this information, the relative mass of each element may be determined.
3) In order to provide uniformity, the atomic mass scale was established
4) The current reference standard is the isotope carbon-12 (
12
12
𝑒), it is assigned a mass of exactly 12
atomic mass units (amu or u). 1 nuclide of
𝑒 = 12.000 000 000.....amu
5) The atomic mass unit of the atomic mass scale it is defined as 1/12th the mass of the carbon-12
nuclide. It is abbreviated as amu or u
12
6) 1 amu = 1/12 (
𝑒) nuclide
7) In looking at masses of the elements on the periodic table it is evident that most values aren’t listed
as whole numbers. Why?
a) These values are actually averages that represent all the naturally occurring isotopes and
the relative abundance in which they are found in nature may be determined by calculation
from any analysis yielding the actual masses of the isotopes and some relative proportion
of their presence (an actual number of nuclides or as percentages of nuclides) A “rough
estimate” may be obtained if the mass numbers are utilized.
Mass Spectrometry:
1) The mass spectrometer or mass spectrograph is a useful instrument in determining chemical
analysis or analysis of isotopes
2) The atoms of material are ionized giving them a positive charge
a) Their charge is usually +1 but may sometimes be greater
3) The ions travel through the magnetic field assuming a cured path which directs them onto a
detector or film plate
a) Number of hits at a particular spot gives the abundance (concentration)
b) ** placement of hits leads to the calculation of mass to charge and eventually identity
4) Different ions are deflected by the magnetic field by different amounts. The amount of deflection
depends on: the mass of the ion. Lighter ions are deflected more than heavier ones.
5) The charge on the ion- ions with 2 (or more) positive charges are deflected more than those with 1
positive charge
6) These two factors combined into the mass/charge ratio. Mass charge ratio given the symbol m/z or
m/e
a) For example if an ion had a mass of 28 and a charge of +1, its mass/charge ratio. It makes it
simpler to talk about this if we assume that the charge on all ions is it. Most of the ions
passing through the mass spectrometer will have a charge of 1+ so that the mass/charge
ratio will be the same as the mass of the ion
7) Assuming 1+ ions stream A has the lightest ions stream B the next lightest and stream c the
heaviest. Lighter ions are going to be more deflected than heavy ones
10/2018 Chemistry Unit 1: Fundamental Properties of Matter
The questions that chemistry attempts to answer:
1) What is this matter I have?
2) How much do I have?
3) How can I change it?
4) How much can I get and how fast
Chemistry - A branch of science dealing with matter, its properties, and its changes
Organic Chemistry - Carbon Containing compounds, this branch of chemistry used to deal with living
organisms, but has broadened to human made compounds
Exceptions 1) oxides (CO and CO2)
2) carbonate (CO3^-2)
Inorganic Chemistry- Primarily focuses on non-carbon containing compounds (oxides and carbonates)
Biochemistry- structure, composition, and reactions within living organisms
- Bioinorganic, biorganic, biophysical
Physical Chemistry- study of matter and its behavior on the molecular and atomic level, how chemical
reactions occur and how complex structures are formed
Analytical Chemistry- Studying matter in regards to its composition and amount qualitative and
quantitative analysis
Scientific Method:
1) Problem Identification: based on observation, olfactory, auditory, visual, tactile, gustatory senses
a) Either qualitative or quantitative
b) Empirical evidence
c) Often a question is developed
2) Generate a Hypothesis:
a) A tentative explanation may involve research of literature often in the “if...then” format
provable hypothesis
3) Experimentation:
a) Tests the hypothesis
b) Component of experimentation- variables, controls/control groups, constants, results/data,
i)
Independent variable: variable that is changed
ii)
Dependant variable: measured/observed variables
iii)
Controls/control group: point of comparison data results
4) Conclusions
a) Conclusions are developed after results/data is evaluated frequently data is looked at to
identify a pattern or regulatory evaluation of facts is often done by making inferences
i)
inferences/interpretations of observations
ii)
Inferences often utilize our prior knowledge/experience
iii)
Inferences must be validated
Progression Towards Certainty:
1) Hypothesis: limited supporting data, based primarily on initial observation and research. May be
disproven or modified as data indicates
2) Theory: often takes the form of a model
a) Results when a hypothesis is tested for a substantial amount of time with enough
supporting data
b) Provides explanation
c) Based upon interpretations
d) Capable of change with new information
3) Theory and Models
a) Models may represent real/hypothetical concepts
i)
Real objects represented by “scaled” forms/drawings
b) Hypothetical models- developed from surrounding/circumstantial data or evidence, models
increase understanding, visualization, enhance communication, predict future event,
explain past circumstances
4) Scientific/Natural Law
a) Law = a summary of observed behavior in nature a statement, at times mathematically
expressed, what is
b) This is in contrast to a theory which attempts to explain why something happens
5) Matter
a) Anything that has volume/mass
6) Chemical Reactions
a) Changes resulting from interactions between matter
Mixtures:
- consist of varying amounts of substances
- Substances: elements or compounds
- Homogeneous = solutions ( same properties throughout)
- Heterogeneous (consists of many different phases)
- Solute: frequently loses phase identity, exists in lesser amount
- Solvent: material which dissolves the solvent
- Solution formation: uniform distribution throughout solvent
- Unsaturated Solution: chemical concentration lower than its equilibrium
solubility
- Saturated Solution: solution containing maximum concentration of solute
- Supersaturated Solution: solution in which the solvent is actually holding
more solute than it should for the conditions, rather unstable
- Solution formation: solutions result when atoms, ions, or molecules which
are on the order of 10^-9 m or smaller in diameter are dissolved in a solent
- Suspensions: when relatively large particles are distributed throughout the solvent,
eventually they fall out, generally have a diameter of 10^-6 m (micrometers)
- Colloidal Suspensions: particles intermediate in size between those in
solutions/suspensions which can be mixed without settling out, range from 10-8 to
10^-6m in size
- Types of Colloidal Dispersions:
-
- Aerosols- solid/liquid particles in gas
- Sols- solid particles in liquid
- Emulsions- liquid dispersed in water
- Foams- gas dispersed in liquid
- Gels- gas dispersed in solid
Tyndall Effect: not observed in solution, colloidal dispersions observed,
suspensions vary
- Always transparent, light passes through w/o scattering, solution is
homogeneous, cannot be filtered
- Liquid colloidal dispersion: intermediate, light will be reflected, cannot be
filtered
- Liquid suspension summary: cloudy and heterogeneous, particles larger
than 10^+4 angstroms
- Allowed for filtration
- Mixture - physical combination separated by physical means
-
-
Filtration: used to separate solid/mixtures
- Sometimes a Buchner funnel will be used allows them to
apply a vacuum during filtration
- Sidearm flask
- Filtrate- liquid component
- Crystallization: Process used to remove a solid solute from a solution, crystals form
when solute reaches a maximum -- saturation
- Concentration - expression of an amount of one substance present
throughout another
- cooling/evaporation will trigger the crystallization process. A reduction in
temperature can trigger crystal formation because solubility decreases with
temperature
- evaporation - increases concentration by removing solvent rate of evaporation
influences crystal size, faster rate, smaller crystals
- Chromatography: a method of separating a mixture of chemicals in gs/liquid form,
by letting them crawl slowly along another substance, surface
- Stationary phase: substance providing the surface along which the mixture can
crawl/separate
- Mobile phase: drags the particle across the surface of the stationary phase
- Results from surface attraction, each attracted at a varying degree
- Those with string attraction will be retained in their spot for an
increased time
- Adsorption- surface attraction
- Objective is to create as much separation between components of a mixture
as possible. Therefore stationary phase requires substantial surface area
- Paper Chromatography: a spot of the mixture is placed near one end of some filter
paper- climbs due to capillarity
- The solvent strikes and dissolves the mixture carrying it along the paper
- Separates during travel due to varying rates of attraction
- At times reagents, developers must be applied
Rf and identifying components: Rf= Retention factor, ratio of distance moved by mixture to
distance moved by solvent (solvent front ) both from common origin
Compare Rf values of unknown and known
Rf= (distance to center of component)/(distance moved by solvent front)
TLC (thin layer chromatography): almost identical to paper chromatography, but
glass/plastic coated with absorbent material is used instead of paper
-
-
-
Column Chromatography: used to separate mixtures glass tube/jar containing highly
absorbent solid silica, silica gel, alumina, or solid coated with liquid
- Eluent poured into column and moves down due to gravity
- Due to different attraction to the adsorbent surface, moves at different rates
GC Unit (Gas Chromatography)- carrier gas (frequently either helium or nitrogen)- liquid
component becomes gas, forced through column, separating the components
HPLC (High Pressure Liquid Chromatography)- instrumental chromatography in which a
solution is placed under high pressure, pushes through a packed column, identifying and
separating the materials
Distillation- used to separate a mixture of liquids, must have distinctly different boiling
points
A liquid is converted to apor which leaves the mixture, vapor travels through a condenser
which turns vapor to liquid, and is then caught in a vessel
- If separating a mixture of liquids, each component comes off near its boiling point
- Generally requires multiple distillations
Substances:
Implies that a substance is pure, even though in reality, we seldom find a strictly pure substance
Classification further delineated to
- Compounds: 2 or more types of atoms
- Represented by chemical formula, consisting of chemical symbols and subscript, where 2
or more elements chemically bond
- Molecules are the smallest representative unit of compounds
- Covalently bonded!
- Ex. NH3
(ammonia)
-
-
Ionic compounds do not form molecules- but tend to form large crystals
- Ionic compounds form as a result of ionic bonds
- Occur when one or more electrons is transferred from a metal to a nonmetal
- Bond between oppositely charged ions- neutral collectively
- Because ionic compounds tend to form crystals composed of an unknown, but vast
number of ions, the smallest representative unit is a formula unit
- Formula unit represents the simplest ratio between positive and negative
ions (aka empirical formula)
- Compounds are capable of being simplified by chemical means, such as heating, or
electricity (electrolysis)
- Fyi: water is a poor electrical conductor alone
Elements: One type of atom
- Represented by the atom- which exhibits all properties of an element
- May be found in free element state - monatomic in nature
- Ex. He and Ar
- Some occur diatomically
- Ex. H2, N2, O2, etc
- Some tetratomic or octatomic
- Allotrope- different forms of the same element
- carbon - graphite, diamond
- Oxygen- O2 and O3 (Ozone)
- Represented by a chemical symbol
Classification of Matter (By Phase):
- Based on packing, motion, and shape
- Solids: definite shape and volume
- Crystalline solids: within a lattice by strong interparticle attractions
- Movement is very limited, primarily vibrated, less than 1% free space
- Amorphous solids: don’t have a particle arrangement of particles
- Ex. waxes, rubbers, plastics
- Liquids: definite volume, indefinite shape
- Interparticle attraction moderate
- Struggle to leave body of fluid
- Liquids considered fluid
- Fluid- substance capable of flowing
- Pascal’s Law: is a principle in fluid mechanics that states that a pressure change
occurring anywhere in a confined incompressible fluid is transmitted throughout
the fluid such that the same change occurs everywhere.
- gases : neither definite shape nor definite volume
- Weak interparticle attraction
- Constant, random motion
- Because of limited attraction, has large range of motion
- Fluid
- Large amount of free space (~90%), and as a result, are easily compressed
Properties of Matter:
2 types: chemical or Physical
- physical : anything observed without altering chemical identity or composition
- Observations may consist of quantitative or qualitative measurements or traits
- Melting point
- Boiling point
- Electrical conductivity
- Thermal conductivity
- Color
- Odor
- Hardness
- Ductility
- Malleability
- Density
- Water has a maximum density of 1.000 g/mL at 3.98 degrees C, used to
describe spec. Gravity
- Chemical: generally involves destruction of the matter in order to identify
- Identified as one substance is chemically transformed into another
- Typically involves interaction with other forms of matter
- Ex. iron has the ability to rust, forms ferric oxide
- A set of chemical properties is more challenging to distinguish
- Flammability
- Combustibility
- Reaction with chlorine
- Oxidation
- Intensive: independent of of sample size
- Extensive: relies on sample size
Physical Changes:
- Boiling vs. evaporation
- Evaporation is a surface process, boiling occurs within the body
Chemical Changes: Results in the formation of one or more new substances
with a unique set of properties
- Precipitate formation
- Color change
- Energy change
- Exothermic
- Endothermic
- Evolution of a gas
Temperature:
Intensive property of matter
A measure of the average kinetic energy of the particles making up a substances
- Kinetic energy: energy associated with the movement of matter (KE=y^2mv^2)
- Any temperature change must be a result of changes in the velocity of the particles of matter
- Thermometer- graduated glass tube with bulb attached, center of tube is hollow and evacuated as
liquid is warmed, it expands, rising up the evacuated tube, as liquid cools, it contracts
- 3 temps scales: Fahrenheit, Celsius, kelvin
Honors Chemistry Unit 1 Part II:
Scientific Notation (Exponential Notation):
Lower numbers of significant figures
General form is M x 10𝑒 × 109- n is any positive of negative integer
Scientific notation has the decimal placed in standard form- to the right of the first nonzero digit
Examples of condensing:
0.000 000 000 753 kg is
7.53 × 10−10 kg
123,000,000,000 pL is
1.23 × 1011 pL
10,00 g
1 × 104 g
Examples of expansions:
5.14 × 105 mm
= 514,000 mm
3.45 × 10−4kg
=0.000 345kg
Recognize, if the power of ten is positive, the number’s value must be greater than 1, if the power of ten is
negative, the number’s value must be less than 1
Shifting the decimal to the left increases the power of ten in a positive manner
3.4 × 103ft = 0.034 × 105 ft
Shifting the decimal to the right increases the power of ten in a negative number
3.4 × 103ft = 34 × 103ft
Shifting the decimal to the left increases the power of ten in a positive manner
3.4 × 103ft = 3,400 ft
Shifting the decimal to the right increases the power of ten in a negative manner
3.4 × 103ft = 0.003 4ft
Adding in Scientific Notation:
All terms must have the same power of ten, if any terms have different powers of ten, temporarily shift
the decimal to obtain like powers (Lines up the decimal)
Once all terms have the same power of ten, add the coefficient terms, round off for significant figures,
retain the power of ten
If necessary, put in STANDARD FORM!!
Example:
E.g. 2.3 × 105 𝑒𝑒 + 4.55 × 103 𝑒𝑒
230 × 103 𝑒𝑒 + 4.55 × 103 𝑒𝑒 = 234.55 × 103 𝑒𝑒
Proper form: 2.3 × 105 𝑒𝑒
E.g. 3.4 × 10−2 𝑒𝑒 + 8.2 × 10−2 𝑒1𝑒 = 11. 6−2 × 10−2 𝑒𝑒
Proper form: 1.16 × 10−1 𝑒𝑒
Subtracting in Scientific Notation:
All terms must have the same power of ten, if they different powers, temporarily shift the decimal
Once all the terms have the same power of ten, subtract the coefficient terms, round off for significant
figures, retain the power of ten. IF NECESSARY, put in standard form.
Example:
E.g. 8.1 × 103 𝑒 − 3.3 × 103 𝑒 = 4.8 × 103 𝑒
E.g. 9.11 × 10−4𝑒𝑒 − 7.21 × 10−5 𝑒𝑒
9.11 × 10−5 𝑒𝑒 − 7.21 × 10−5 𝑒𝑒 =
83.89 × 10−5 𝑒𝑒
83.9 × 10−5 𝑒𝑒
−4
Proper form: 8.39 × 10 𝑒𝑒
Multiplication in Scientific Notation:
Multiply coefficient terms of all factors
Sum the powers of ten for all factors
Round for significant figures as needed
Correct for standard scientific notation form[
Example:
E.g. 1.23 × 103 𝑒 × 7.3 × 102 𝑒
= 8.97 × 105 𝑒2
Proper form:9.0 × 105 𝑒
Division in Scientific Notation:
Divide the numerator of dividend coefficient by the denominator or divisor coefficient
Subtract the power of ten of the denominator of divisor from the power of ten of the numerator of
dividend
Round off for significant figures
Correct for standard scientific notation form
Example:
E.g. 4.5 × 10−3 𝑒𝑒 ÷ 5.22 × 10−2 𝑒𝑒
= 0.862 069 × 10−1 (note, NO UNITS)
= 0.86 × 10−1
Proper form: = 8.6 × 10−2
Scientific Notation and Powers of Exponentials:
The digit term is raised to the indicated power and the exponent is multiplied by the number that
indicates the power
E.g. (2,4 × 104 𝑒)3
=(2.4)3 × 10(4×3) 𝑒3
= 13.824 × 1012 𝑒3
Proper form: 1.4 × 1013 𝑒3
When to Use Scientific Notation this Year
I am anticipating scientific notation to be used for numbers with a value of 10,000 or larger, of for
numbers with a value or 0.000 1 or smaller
Also, scientific notation is just another means of expressing a numerical value, therefore significance
cannot change when moving into or out of standard form
Uncertainty:
All measurements in science have some degree or uncertainty
Every measuring instrument/tool has a limitation - none measure to infinity
Each technician/operator has a various skill level
Accuracy- Accuracy is the closeness of a measurement or value to the real, true, or ACCEPTED value
It is conveyed through the use of the term ERROR
Absolute errorThe difference between the experimental(measured) value and the accepted value
𝑒𝑒
Relative Error
The same as percentage error
Provides an idea or the magnitude of impact or the error as it compares the absolute error
to the
accepted value
Precision:
The closeness of a group of measurements made in a similar manner
It is conveyed by the term DEVIATION
May also refer to the exactness to which an instrument may measure
Absolute deviationthe difference between an experimental value and the mean of a set of values
Good scientific work would imply making/taking multiple values
𝑒𝑒 = |𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒 𝑒𝑒𝑒𝑒𝑒 − 𝑒𝑒𝑒𝑒 𝑒𝑒 𝑒𝑒𝑒|
Relative deviationThe relative deviation compares the average deviation of the set of data to the
magnitude of the mean; it is expressed as a percentage
𝑒 𝑒 = (𝑒 𝑒(𝑒𝑒𝑒) /𝑒) × 100
Significant Figures:
Significant figures are important in science as they convey uncertainty in measurements and calculations
involving measurement.
In science, there are no repeating decimal values. Every value has a specific number of digits which are
used to express it due to uncertainty in measurement and measuring instruments.
2 different types of numbers:
- Exact
- Measured
Infinite Significance:
-
Exact numbers have infinite significance
- These values are obtained by definition or counting
- For example, the number of students in a class is 21...there aren’t 20.5 nor 21.35.
You have a student or you do not
- NO uncertainty
-
-
By definition: likewise, a centimeter is defined as 100th of a meter
- 1 cm = 10−2NO uncertainty
- No measuring tool employed in these determinations
Measured numbers
-
Determined with a measuring device so these values have uncertainty
- Every measuring tool has its own limitation
- Each also has a degree of uncertainty built into its manufacture
- The operator also adds a degree of uncertainty
- EXPRESSING MEASUREMENT:
- An estimate is needed for one units place beyond the finest
graduation
- Composition of a significant figure: all but one significant figures are
known with certainty
- The last significant figure is only the best possible estimate
- Recognize, an estimated value still conveys meaning
- However, the “implied uncertainty” is always +/- 1 of the last
significant digit in the value
Conveying Uncertainty:
-
-
In order to communicate reliability of a measurement or data, the uncertainty in the measurement
must be expressed
- Provides additional meaning to the piece of data
- Significant Figures: a simple means for conveying uncertainty associated with a
measurement
- They are frequently employed when only one or two measurements are made
instead of a series of measurements from which an average (mean) and an actual
uncertainty could be determined
- What is a significant figure/digit? One known with “reasonable reliability”
- What digits are significant? All digits in a measurement are considered
significant except for place holding zeros whose function is to provide the
value with its magnitude
Therefore, magnitude of a number has nothing to do with its
significance
Rules for Determining Significant Digits:
1) All nonzero digits (1-9) are significant
2) All zeros between significant digits are significant
3) Zeros ending a number to the right of the decimal are significant
4) Zeros ending a number to the left of the decimal point are not significant unless indicated
a) Indications may be expression of the decimal point or a repetend bar over the last
significant zero
5) Zeros to the left of significant digits are not significant
Calculators:
When using calculators a correct answer must be determined; recognize calculators only do what they are
told; and don’t know the correct answer
When you use your calculator your answer can only be as precise as your worst measurement
- MAKE NOTE!! All problems given on tests and quizzes in this class will be expressed to the desired
significance. Treat them accordingly.
Operations Using Significant Figures:
Addition and subtraction-
1) Identify the least significant value on a units/decimal place basis
a) Value with the largest uncertainty (+/-)
2) Perform the operation (add/subtract)
3) Round the result to the same units/decimal place as the least significant value
i)
2.45g + 13g + 0.88g = 16.33 g →
ii)
(+/-0.01g) (+/-1g) (+/-0.01g) →
iii)
Therefore, the value 13g has the greatest uncertainty, and therefore will cause the
result to be rounded to the 1’s of a gram
iv)
SO, 16.33 becomes 16g for the final answer!
Multiplication and Division
1) Identify the least significant value based upon the number of significant digits
a) Based upon how many sig. Figs are present
2) Proceed with operation (multiply/divide)
3) Round the result to have the same number of significant digits as the least significant digits as the
least significant value
i)
374.0 𝑒𝑒2 / 22.0cm = 17cm →
ii)
(4 sig. figs) (3 sig. figs) →
iii)
Therefore, the answer must have 3 significant figures
iv)
Answer is 17.0 cm
(1) 2.443 cm x 8.04 cm x 4.0 cm = 78.578 606 4𝑒𝑒3
(2) (4 sig figs) (5 sig figs) (2 sig figs)
(3) Answer will be 79cm^3
Mixed Operations:
-
Significant figures must be accounted for every time that the operation changes!!! Follow order of
operations
Rules for Rounding:
1) If the first digit dropped is less than 5, leave the preceding number unchanged (3.133 becomes
3.13)
2) If the first digit dropped is greater than 5, increase the preceding digit by 1 (3.127 becomes 3.13)
3) If the first digit dropped is exactly 5, round off to make the preceding digit to an even number
Scientific Notation and Sig. Figs:
-
-
The coefficient of the term must contain only and all significant figures
- 2,300 in = 2.3 x 103 in
- 0.245 00 Mm = 2.450 0 x 10−1Mm
Expanding w/ sig figs
- 2.40 x 104 g = 24,Ō00
- First zero w/ REPITAND to indicate significance
Measurement:
-
A quantity is represented by a number and a unit
The unit in a quantity and hence science is extremely important. It provides the standard to which
our measurement is compared
- UNITS provide MEANING
SI System:
-
The SI system is based primarily upon the metric system
The SI system employs prefixes in order to express quantities reasonably based upon their
magnitude
Prefixes with values greater than one
Need to memorize!!
Bare in mind case!
Prefixes with
values less than
one
ALSO memorize!
Length:
-
Distance covered by a line segment connecting two points
SI standard unit is the meter, symbol ‘m’
English to metric conversion
- [1 inch = 2.54 cm]
- [1m = 39.37 inches) ← MEMORIZE
Volume (capacity):
-
-
Amount of space defined by 3 dimensions
SI standard unit is the 𝑒3
The liter is the older metric unit of volume that is frequently employed today
- L is symbol
- Useful equalities: 1L = 1𝑒𝑒3 and 1mL = 1𝑒𝑒3
English to metric equivalent
1.056 688 quarts = 1L; L= 1.057; 1qt = 946mL
Equations for Volume!
Memorize:
-
1 bushel
- 4 pecks
1 peck
- 8 dry quarts
1 dry quart
- 2 dry pints
Liquid Measure:
-
1 gallon = 4 quarts
1 quart = 2 pints
1 pint = 2 cups
1 cup = 8 floz
Units of Length Cubed:
- 1 𝑒𝑒𝑒𝑒3
- 1 𝑒𝑒𝑒𝑒3
- 1 𝑒𝑒𝑒𝑒= 128𝑒𝑒3
Area:
-
Measurement of a surface as defined by 2 dimensions
SI standard is the 𝑒2
1𝑒𝑒2 = 640 acre
-
1 acre
1𝑒𝑒2
1𝑒𝑒2
Mass:
Measurement of a quantity of matter
- Matter is anything having mass and volume or anything exhibiting inertia
- The property of matter which resists change in state of motion
- EITHER at rest, or in motion
- At constant velocity, an object at rest stays at rest, an object in motion stays in
motion, thus speed and direction REMAINS THE SAME
- VELOCITY is a VECTOR quantity
- Unless acted on by another force
- MASS is a measure of inertia
- Kick a soccer ball vs kick a
bowling ball
- Newton’s second
law:
- Standard SI unit is the kilogram,
kg
- The only prefixed SI unit
Weight: the measurement of force acting on a mass
- Common field of force is gravity
- Force of attraction between any
two masses
- Gravity is a weak force dependant
upon
- MASS & DISTANCE
-
SI standard unit of weight is the newton (N)= 1kgm/𝑒2
English (US)
- Poundal - a unit of force equal to that required to give a mass of one pound an
acceleration of one foot per second
- Pounds-force
-
The pound-force is equal to a mass of one avoirdupois pound multiplied by the
standard acceleration due to gravity on Earth, which is defined as exactly 9.80665
meter per second squared
- 1 pound force = 4.448 N
- 1 pound (Lb.) = 4.448 N
Mass:
Why is mass preferred over wight?
- Independent of gravity
English to metric equality
1 Lb. = 453.592 37 grams
USE: 1 Lb. = 453.6 g
- Measured by a balance
Time:
Measured interval between two occurrences or events
SI standard unit is the second
Derived Units:
Fundamental or base units are obtained from making one measurement.
If two or more measurements and hence their units are combined, the resulting unit is referred to as a
derived unit.
e.g. m3 , g/cm3, ft. /sec , Lbs./in2
Dimensional Analysis:
Dimensional analysis or factor label is a problem solving method that places an emphasis upon the
accountability of units
One of its primary functions is unit conversions
The process involves multiplying a “given” by one or more conversion factors in order to change the units
of the quantity without changing the magnitude of the quantity
Conversion Factors:
Conversion factors are ratios (fractions) with a value of one (1)
These factors may be obtained from equalities or from other “given” pieces of information
This process enables one to use COMMON pieces of information in the problem solving process
Performing Dimensional Analysis:
Start with the “ given” upon which the “find” will depend
This “given” is often placed over one (1) to emphasize its numerator position
Logically, multiply this “given” by ordered conversion factor(s) that will cancel out
units and introduce desired units
Continue multiplying by conversion factor(s) until the final desired unit is obtained
undesired
How to use SI prefixes:
Recommendation:
Place the coefficient of one in front of the prefixed unit, set an equals sign, the power of ten
representing the prefix, and then the base unit.
1 prefixed unit = power of ten base unit
1 kilometer = 10+3 m
When converting between two SI units, the conversion may be performed in:
1)Two steps by converting to the base and then to the desired unit
2) In one step by applying a single conversion factor between the units
Example:
5 mg to kg
k - 10+3
Δ = 10+/-6
m – 10-3
10+3 – (-3) = 10+6 10-3 – (+3) = 10-6
Chemistry Unit II:
Periodic Table: One of the most important tools of the chemist is the periodic table of the elements. It is
used for element identification, element masses, prediction of properties such as ionization energy, etc.
Johann Dobereiner: In 1829, he classified some elements into groups of three, which he called triads. The
elements in a triad had similar chemical propertie and orderly physical properties. (ex. Cl, Br, I, and Ca,
Sr, Ba)
John Newlands: In 1863, he suggested that the elements be arranged in “octaves”because he noticed (after
arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th
element
➢ Law of Octaves
Newland’s claim to see a repeating pattern was met with savage ridicule on its announcement. His
classification of the elements, he was told, was arbitrary as putting them in alphabetical order and his
paper was rejected for publication by the Chemical Society.
➢ The law of Octaves failed beyond the element Calcium
Dmitri Mendeleev: In 1869 he published a paper organizing the elements by atomic mass, at the same
time Lothar Meyer published his own table of the elements organized by increasing atomic mass.
➢ Competitive with Lothar Meyer: arranged elements in order of properties. It appeared that
repetition of properties was based upon atomic mass. (Initial periodic law)
➢ It is Mendeleev called the father of the modern periodic table, and not meyer. Why?
○ Filed for publication prior to Meyer
➢ One problem with the table… Both had spots where unknown elements should fit. Mendeleev used
his table to determine physical properties of missing elements yet to be discovered.
➢ After the discovery of elements between
1874 and
1885, and the fact that Mendeleev’s
predictions
for Sc, Ga, and Ge were amazingly close
to the actual
values, his table was generally accepted.
➢ However, in spite of Mendeleev’s great
accomplishment, problems arose when
new elements
were discovered and more accurate
atomic
weights determined. Some elements
appeared to
not fit the pattern. (ex. Ar and K,
Co and Ni, te and I, Th and Pa)
○ Mendeleev believed that if the
atomic
weight did not fit into the pattern,
the weights
were incorrect.
Henry Moseley: In 1913, through his work with x-rays, he determined the actual nuclear charge (atomic
number) of the elements. He rearranged the elements on the periodic table in order of increasing atomic
number.
➢ “There is in the atom a fundamental quantity which increases by regular steps as we pass from
each element to the next. This quantity can only be the charge on the central positive nucleus.”
➢ Anecdote: His research was halted when the British government sent him to serve as a foot soldier
in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of
this loss, the British government later restricted its scientists to noncombatant duties during
WWII.
Modified Periodic Law: The properties of the elements are a periodic function of their increasing
atomic numbers.
➢ Periodic: repeats at regular intervals
Included on the Periodic Table: The periodic table is a comprehensive listing of the known and accepted
chemical elements. Most tables included the elemental name, the chemical symbol, the atomic number,
and the average atomic mass.
➢ Reminder: average atomic mass is calculated using naturally occurring isotopes!
Groups: A group on the periodic table is a vertical column of elements. Sometimes groups are referred to
as families.
➢ The reference to a family results from the similarity of properties shared by the elements found in
a vertical column.
➢ Eg. Alkali metals – lower densities than other metals; low ionization energies; largest atomic
radius of their horizontal row of elements; low electronegativities; form positive one cations; soft;
shiny; they all react with water, …….. (1st column)
Periods: Horizontal rows of elements
➢ The elements of a period have the same electron valence shell, but have limited similarity in
properties.
○ Valence- outermost energy level with at least 1 electron
○ There are 7 periods
“A” Groups: “A” groups are collectively known as the representative elements or main block/group
elements.
Some of the more important families or groups:
Group 1/Family IA – alkali metal family
Group 2/Family IIA – alkaline earth metal family
Group 16/Family VIA – chalcogen family
Group 17/ Family VIIA – halogen family
Group 18/Family VIIIA – noble gas family
“B” Groups: the transition metals
Broken out below the transition elements are the inner transition elements.
Elements 58 – 71 are called the lanthanide series
Elements 90 – 103 are called the actinide series
Metals: Metals are elements located left of the staircase on the periodic table, excepting hydrogen (& He if
present)
Metallic properties:
Luster, Conductivity (thermal and electrical), Malleable, Ductile, Generally higher melting points and
densities
➢ Malleable- pounded into a shape, Ductility- ability to be extruded into a fine wire/strand
Nonmetals: Nonmetals are elements listed to the right of the staircase on the periodic table
Nonmetallic properties:
➢ Lower melting points and densities
○ Because, most of them are gasses
➢ Solids are brittle
➢ Insulators- retard or impede the flow of electrical/thermal energy
➢ Dull appearance
Metalloids: Elements which are adjacent to the staircase on the periodic table
★ The exception is aluminum- a metal, occasionally polonium is excluded
➢ These elements generally share some characteristics of metals and some nonmetals (thusly,
transition elements)
➢ May be shiny or dull
➢ May conduct but not well as a metal would
○ Hence, called semiconductors
○ In technology, Silicon tends to be the most important, generates a tremendous
amount of heat, and thus require a fan to cool
Alkali Metals: Highly reactive metallic elements in Group 1 or Family IA- most reactive
➢ react with water to form hydrogen and alkaline solutions (bases); burn in air
○ al-quili means wood ashes
○ -term dates back to ancient times; people discovered that wood ashes mix with
water to produce slippery solutions that can remove grease
➢ one outer electron; by losing this electron they become a +1 cation
➢ Soft, but can be cut with a knife
➢ Shiny, but dull quickly due to oxygen and water in air
➢ Good conductors; Gaseous states at high temperatures become plasmas
○ e.g.Li, Na, K, Rb, Cs
Alkaline Earth Metals: Elements found in Group 2 (family 2a) - second most reactive
➢ Oxides are not very soluble and are found in the earth, from which comes the idea of the
name “Earth”
➢ Oxides unable to light on fire- already possess full complement of oxygens
➢ Reactive metallic elements with 2 electrons in the outermost energy level
○ Results in the formation of a +2 cation
➢ Compared to alkali metals: harder, denser, stronger, higher melting points, lower reactivity
than alkali metals
➢ Eg. Be, Mg, Ca, Sr, Ba, Ra
Chalcogens: Elements found in Group 16 (Family VIA) of the table
➢ The name means “ore former” (greek) as many ores contain sulfur or oxygen- chalcos - ore
and gen - forming
○ Eg. cinnabar (mercuric sulfide); galena (plumbous sulfide); bauxite (hydrated
aluminum oxide); hematite (ferric oxide)
➢ These elements have 8 electrons in their outermost energy level
○ Tend to acquire 2 electrons, thus forming a -2 anion
Halogens: Nonmetallic elements in Group 17 (Family VII A), that have 7 valence electrons and combine
with metals to form salts - chemically the most active nonmetallic family
➢ Tend to capture electron forming anion -1
○ Most active element - fluorine
➢ Term “halogen” comes from Greek meaning “salt forming”
➢ Contains fluorine- the most chemically active element
➢ Exist diatomically
○ e.g. F2, Cl2, Br2, I2, At2
Noble Gases (Prior name inert gases): Elements in Group 18 (Family VIIIA) that are characterized by low
reactivity
➢ The term comes from “noble” people, did not associate with anyone other than their kind
➢ Characterized by one octet (8) electrons in their valence shell; very stable
○ Exception of helium- “duet”, but this still completes its valence shell
○
○
Colorless, odorless, very stable (unreactive)
e.g. He, Ne, Ar, Kr, Xe, Rn - Most concerned with radon, due to its radioactivity
Transition Metals: Elements located primarily in the shorter groups
➢ These elements share metallic properties and character, however, to varying degrees
➢ These elements frequently have more than one oxidation state (form more than one
positive charge)
➢ Many of these metals have high melting points
➢ Frequently contain compounds which have a color to them other than white
○ Eg. copper compounds may be green or blue
Inner Transition Metals: Once referred to as the rare earth elements
➢ This block made up of two rows of elements: one the lanthanide series, the other the
actinide series
○ Now can be produced synthetically
➢ These elements broken out of table for convenience; to make it shorter and therefore; fit in
a page
➢ Many of these elements are white or silvery in color, they tend to be good conductors, the
elements are lustrous, but tend to tarnish quickly in air
➢ Lanthanides: Shiny metallic transition metals (58-71)
➢ Actinides: Shiny metallic transition metals, many of these elements are synthetic or manmade, tend radioactive (90-103)
Hydrogen: The most common element in the universe
➢ It is considered a family all of it own due to its unique properties and characteristics
○ Physically has properties like a nonmetal, but chemically behaves more like a metal
➢ Composed of one proton and one electron making it the simplest atomic structure
○ Protium, deuterium, tritium
➢ Hydrogen reacts with numerous elements
➢ Found elementally in its diatomic condition, H2
Representative Elements: The representative or main block elements are those with an “A” heading- the
longer groups (gps, 1, 2, 13 thru 18)
➢ Many of these elements tend to be rather reactive (possess; halogens, alkali metals,
alkaline earth metals, chalcogens)
○ Excepting the noble gases
Chemical Formulas: Are a means of describing the composition of a chemical compound and the
elements present in that compound
➢ Composed of chemical symbols and subscripts
➢ Free Elements: Not combined with another element in a compound
○ Ex. Au (gold), Cu (copper), and He (helium)
○ Free elements are frequently represented by a chemical symbol
➢ However, most nonmetallic elements excepting the noble gases exist elementally in
molecules composed of more than one atom
○ P4, S8, O3
➢ Chemical formulas specify the composition of a substance, if the substance is a chemical
compound
○ e.g. Fe2O3 (ferric oxide or iron(III) oxide)is composed of the elements iron and
oxygen in a 2:3 ratio
○
e.g. CO(NH2)2 {urea} expands to CON2H4, but parentheses often group atoms to
show the compound’s actual structure
Hydrates (Hydrated Compounds): Hydrates are compounds that contain water molecules.
➢ The water molecules are usually bonded weakly into the compounds structure
➢ Many of these compounds are ionic in nature
○ For example plaster: CaSO4 • 2 H2O
○ For example epsom salts: MgSO4 • 7 H2O
➢ The formula of a hydrate is written as the anhydrous formula (or formula of the “salt”)
followed by a raised dot (•), a coefficient (except “1”), and then the formula for water (H2O)
○ Na2CO3 • 10 H2O
➢ The names of hydrates are expressed as the name of the anhydrous compound followed by
a prefix attached to the word hydrate.
○ sodium carbonate decahydrate
○ Prefixes: Prefixes: 1 is mono- 2 is di- 3 is tri- 4 is tetra- 5 is penta- 6 is hexa7 is hepta- 8 is octa- 9 is nona- 10 is deca➢ The dot indicates the water molecules are “weakly “ attached to the anhydrous crystal
○ Therefore, they usually can be removed by gentle heating
➢ A coefficient of one is NEVER EXPRESSED!!
➢ Also, the anhydrous compound is generally an ionic compound or “salt”
Law of Definite and Multiple Proportions:
➢ Definite:Atoms combine in fixed, small whole number ratios.
○ e.g. 1 to 1, 1 to 2, 2 to 3, etc.
This coincides with Joseph Proust’s Law of Definite Proportions which is also referred to as the
Law of Constant Composition. The law states that; compounds are composed of elements
combined in a fixed ratio by mass.
○ e.g. Always find 2.02 g H to 16.00 g O in water or 12.01 g of C to 32.00 g O in carbon dioxide.
➢ Multiple: Dalton developed a law known as the Law of Multiple Proportions which states; that the
ratio of a fixed mass of one element when combined with the multiple masses of another element
may be expressed in small whole numbers.
○ In simpler terms, the same element can unite in more than one way with another element
to form several compounds.
Chemical Equations:
➢ Represent chemical reactions
○ 2 HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(l) +CO2(g)
➢ HCl and CaCO3 are called reactants
➢ CaCl2, H2O,CO2 are called the products
➢ Reactants are separated from products with “” that means “yields” or “produces”
○ Symbols of Substances; States matter: for solids use (s), liquids (l), gases (g), vapor (v),
crystalline solids (c), and for substances dissolved in water (aqueous solutions) use (aq).
Balancing Equations:
➢ Coefficients- numbers in front of chemical formulas/chemical symbols to indicate the number of
molecules/atoms of each type
➢ Balancing is achieved by adjusting coefficients. DO NOT USE SUBSCRIPTS!!
➢ Recognize, balancing is performed in order to achieve the Conservation of Atoms
➢ Equations are balanced due to the Law of Conservation of Mass
Molecules Form When Nonmetallic Elements Combine:
➢ Molecules are neutral particles made of 2 or more atoms.
➢ Many molecular compounds contain hydrogen:
Group
Noble
Period IVA
VA
VIA
VIIA
Gas
2
CH4
NH3 H2O
HF
Ne
3
SiH4
PH3
H2S
HCl
Ar
4
GeH4
AsH3
H2Se
HBr
Kr
5
SbH3
H2Te
HI
Xe
Organic Compounds:
➢ Organic compounds are compounds containing the element carbon.
○ Exceptions are: carbonates (CO3-2), CO, and CO2
○ A large number of organic compounds are composed using carbon, hydrogen, oxygen,
additional elements found are nitrogen, phosphorus, and sulfur
Alkanes:
➢ Alkanes are saturated
hydrocarbons (contain only C &
➢ Always have a ratio of atoms
CnH2n+2
➢ Named using a prefix designating
number of C atoms
➢ All have –ane suffix
H)
the
Alkenes:
➢ hydrocarbons with fewer H than the alkanes.
○ Unsaturated hydrocarbons
○ They have the relationship; CnH2n. They possess a double bond.
○ Names use the same prefixes, but have the suffix -ene.
■ Identify the longest carbon chain possible in the molecule with the double bond. It
is named with the corresponding prefix and the suffix – ene.
■ Number the carbons of the chain so the double bond carbons have the lowest
numbers. Place the carbon number and a dash in front of the name.
Alkynes:
➢ Alkynes are compounds possessing a triple bond between two carbon atoms in the molecule.
(unsaturated)
➢ Alkynes have the general formula; CnH2n-2
➢ The simplest alkyne is ethyne or acetylene
○ The next simplest alkyne is propyne
■ Naming alkynes:
Determine the longest carbon chain containing the triple bond
Number the carbons starting at the end closest the triple bond.
Use the proper prefix with the – yne suffix.
Indicate the triple bond position with a number in front of the name indicating the
closest carbon to which the bond is attached
Alcohols:
➢ Are carbon chains in which one H is replaced in an alkane with an -OH group
➢ Naming - Same prefixes, the ane suffix, but the “e” is replaced by
– ol
➢ If the –OH (hydroxyl group) appears on a carbon other than a terminal carbon atom, indicate its
location on the chain using a number referencing the lowest number carbon possible to reference
the hydroxyl group.
Ionic Compounds:
➢ Positively charged ions are called cations
➢ Negatively charged ions are called anions
➢ Subscripts in the formula always specify the smallest whole-number ratio of ions required to make
a neutral combination (formula unit)
○ Due to ionic compounds forming crystals; vast arrays of combined ions
○ Empirical formula: simplest ratio between the combining ions or atoms of a compound
■ Writing Ionic Compound Formula Units:
● The cation written first followed by anion
● The subscripts of the formula must produce an electrically neutral formula
unit
○ Referred to as electroneutrality
● The subscripts should be set to the lowest possible whole number
○ So as to create the empirical formula
● In order to write the chemical formula for an ionic compound, the charges
of the combining ions must be known.
■ Naming Ionic Compounds
● Name the cation first; if the cation is grom group, 1, 2, or 13, merely write
the element’s name
○ B, Al, or Ga or Zn, Cd, Ag
○ If the element is one of the transition metals or a post transition
metal; determine the charge on the metallic ion and write its name
with the “-ic: or “-ous” suffixes or using the stock system (roman
numerals)
○ In order to apply the Stock System or the traditional name to a metal
forming multiple positive ions
■ Determine the total negative charge in the compound based
upon the formula
■ Change the negative sign to a positive sign as all compounds
are neutral
■ Divide the positive charge by the number of positive ions in
the formula (metal ions)
What Ions Form?
➢ Noble gases especially stable, and therefore have low reactivity
➢ Main group electrons will frequently gain or lose electrons to have the same number as the nearest
noble gas (generally an octet in the valence shell)
○ Metals form cations by losing
○ Nonmetals form anions by gaining
■ Fyi, metals tend to lose, nonmetals tend to gain
Properties of Ionic Salts:
➢ Electrical conductivity: While not a conductor in the solid state, it does conduct in the molten or
solvated state.
➢ Melting/boiling points: high melting and boiling points due to strong interionic attractive forces
➢ Structure: Hard, yet, brittle, crystals
➢ Solubility in water: Many, if not all, ionic compounds dissolve in polar solvents such as water
Traditional Names of Elements with Multiple Oxidation Numbers:
➢ Oxidation numbers are values assigned an element in an attempt to describe electron charge
distribution within a compound (molecule or formula unit).
➢ It is usually an indication of the charge on an ion formed by that element.
➢ When naming metals having more than one oxidation number (exhibiting more than one ionic
charge), a suffix of “-ic” is applied to the name of the element for the ion of greater charge and a
suffix of “-ous” is attached for the ion of lesser charge.
○ NOTE: Sometimes the suffix is attached to the name of origin!
The Stock
➢ In
System of Naming:
➢ IUPAC recognized
naming a metal with multiple oxidation
states (having more than one ionic charge), a
Roman numeral is placed in parentheses
immediately following the element’s name.
○
Roman numeral indicates ion’s charge
Naming Elements Having a Negative Oxidation State:
➢ When an element assumes a negative oxidation state, meaning it has acquired a negative charge,
its name is modified to accept an “-ide”
suffix.
Polyatomic Ions:
➢ End in -ate or -ite
➢ Covalently bonded
➢ Others will have the prefix hydrogen, or
bi,
➢ Nonoxyanions end in -ide
Molecular Compounds:
➢ Molecular compounds are a result of covalent bonding found in molecules
○ These compounds may be organic (having their own nomenclature)
○ These compounds are produced of 2 nonmetals
Molecular Formulas:
➢ Molecular formulas represent the actual composition of molecules
○ These molecules may represent compounds, or they may represent elements
➢ Recognize while all compounds have an empirical formula, only molecular compounds have a
molecular formula
➢ Shows the exact number of each element in a molecule
➢ Some elements have two or more distinct allotropes
○ Ex. O2 and O3
➢ A structural formula shows not only composition, but additionally, arrangement
○ While the molecular formula provides the true composition, the empirical formula is the
simplest
Inorganic Compounds:
➢ Classified as either:
○ Oxides: combination of an element with oxygen
○ Acids: Contain hydrogen for donation
■ Bronston-Lowry definition: A Brønsted-Lowry acid is any substance that donates a
proton. A base is any substance that accepts a proton.
■ Lewis: A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair
donor.
○ Bases: most contain cation and hydroxide anions
○ Salts: technically may be formed by a reaction between an acid and a base- positive ion of
base, negative ion of acid
■ Acids and bases neutralize each other’s properties
■ Both will conduct an electrical current
■ Acids will react with active metals forming hydrogen gas
● Hence, acids frequently considered to be corrosive
■ Bases tend to be ‘slippery’
Honors Chemistry Unit III
The Mole:
➢ The mole is a unit used for expressing a quantity of matter
○ It is used in talking about the amount of substance present.
■ What is a substance? Elements or compounds
○ The mole is defined as the quantity of matter composed of 6.022 141 5 ×
1023 particles of that matter
■ Established by x ray diffraction and electrochemistry
■ For the purpose of this class; use 6.022 × 1023
■ The particles referenced here may be atoms (for elements) ions (for charged
particles), molecules (for molecular compounds), or formula units (for ionic
compounds)
○ The IUPAC definition of the mole is; the amount of substance containing as many
“elementary entities” as there are atoms/nuclides in EXACTLY 0.012 kg
12𝑒 (carbon-12)
■ Carbon-12 is the standard for the atomic mass scale
● The mole is abbreviated mol
■ Therefore, hypothetically, we could state:
● 1 mole of bananas contains 6.022 141 5 × 1023 bananas
● 1 mole of nails contains 6.022 141 5 × 1023nails
● 1 mole of electrons contains 6.022 141 5 × 1023electrons
● 1 mole of Mr. Garman’s students contains 6.022 141 5 × 1023 students
○
○
○
○
○
○
○
■ The value, 6.022 141 5 × 1023, is referred to as Avogadro’s number
● It has the symbol𝑒𝑒
○ It is named in honour of Avogadro, not because he
determined the value
○ Avogadro’s Principle: Equal volumes of gasses under the same
conditions of temperature and pressure are composed of the
same number of molecules
■ The magnitude of Avogadro’s number is staggering……
■ How large a number is 6.022 141 5 × 1023 ? What is the
● Green Pea analogy: If Pennsylvania is covered in 4 feet of green peas,
there will be about one quintillion peas - 1018 . If this falls over the
entire continental globe, 4 feet deep; only 1021 peas. One hundred and
fifty planets the size of the Earth, covered completely 4 feet deep in
peas, finally hits close to 6.022 141 5 × 1023
The mole is a unit used for counting or measuring numbers of very small objects
like atoms. It is very difficult to work with these object on an individual basis,
therefore, working with designated groups (multiples) of these infinitesimally small
objects make physical measurement possible and numbers more manageable.
The mole may be more conveniently related to mass for the purpose of measuring
out a specific quantity of substance.
The periodic table has the average atomic mass listed for each element. This mass is
interpreted to be in amu’s (u’s) for individual atoms of each element
However when working with moles, now the number of atoms is extremely large.
Therefore, moles are represented by the gram mass of the element’s gram atomic
mass.
If, for example 1 mole of carbon-12 is composed of 6.022 141 5 × 1023 nuclidesg of
carbon-12 it has a mass of exactly 12.000 000 000 000 0….. g, etc.
■ The value 12.0000 etc is the gram mass of carbon-12 or represents what is
called the gram atomic mass of carbon-12
This same rationale may be applied to any element; one mole of any elements
contains𝑒𝑒 atoms of that element and is equivalent to the atomic mass of the
element expressed in units of grams (element’s gram atomic mass)
■ E.g. 1 atom of aluminum has a mass of 26.981 amus
■ 1 mole of Al atoms (6.022 141 5 × 1023 ) has a mass of 26.981 grams
The same logic may also be applied to compounds as in the following matter:
■ Formula mass- most frequently used when working with ionic compounds
(metallic cation-nonmetallic anion; metallic cation-polyatomic anion;
polyatomic cation - polyatomic anion; or a polyatomic cation - non metallic
anion)
● E.g. 1 formula unit of sodium chloride is 58.44 amu
●
22.99 𝑒𝑒𝑒 𝑒𝑒
=
1 𝑒𝑒𝑒𝑒 𝑒𝑒
35.45 𝑒𝑒𝑒 𝑒𝑒
1 atom Cl x
=
1 𝑒𝑒𝑒𝑒 𝑒𝑒
1 atom Cl x
22.99 amu Na
35.45 amu Cl
○ +___________
○ 58.44 amu NaCl
■ Gram formula mass - When working with one mole of formula units
●
22.99 𝑒 𝑒𝑒
●
1 atom Nax 1 𝑒𝑒𝑒𝑒 𝑒𝑒= 22.99 g Na
●
1 atom Cl x 1 𝑒𝑒𝑒𝑒 𝑒𝑒= 35.45 g Cl
○ + ___________
○ 58.44 g NaCl
35.45 𝑒 𝑒𝑒
■ Molecular mass - used when working with molecular compounds
●
●
12.01 𝑒𝑒𝑒 𝑒
= 22.99 amu C
1 𝑒𝑒𝑒𝑒 𝑒
16.00 𝑒𝑒𝑒 𝑒
2 atoms C x
= 35.45 amu O
1 𝑒𝑒𝑒𝑒 𝑒
1 atom C x
○ +___________
○ 44.01 amu CO2
■ Gram molecular mass- used when working with one mole of molecules
12.01 𝑒𝑒
= 22.99g
1𝑒𝑒
16.00𝑒 𝑒
2 atoms C x 1 𝑒𝑒𝑒𝑒 𝑒= 35.45g O
●
■
1 atom C x
C
○ +___________
○ 44.01 g CO2
○ The terms gram atomic mass, gram formula mass, and gram molecular mass may
be collectively referred to as the molar mass of the substance
■ Molar mass- the mass of one mole of substance expressed in grams
➢ Using the mole in conversions
➢ The concept of the mole allows for conversions between particles and moles, moles and
mass, and particles and mass
➢ The mole and its equivalents allow for the formation of conversion factors (Used in
dimensional analysis!!)
○ 1 mole or 𝑒𝑒
1 mole or molar mass
____ ______ ________ __________
𝑒𝑒
1 mole molar mass 1 mole
Example: moles to mass- (and vice versa)
1) How many grams are present in 0.700 moles of hydrogen peroxide?
a) 0.700 m H2O2 x 34.02gH2O2 = 23.8
_____________
1 mole H2O2
The mole can also be related to a volume of gas, however, only at conditions of STP
STP means standard temperature and pressure
The standard temperature of a gas is 0.0C, 1.000 atm
Equivalents to 1.000 atm ar: 760 mm Hg, 760.0 torr, 14.7 psi, 101.325 kPa, 1.013 25 bar
Molar volume- the molar volume of a gas at STP is 22.4 L
Meaning at conditions of STP, one mole of any gas will occupy 22.4 L of space
Conversion factors formed;
22.4 𝑒 𝑒𝑒𝑒
1 𝑒𝑒𝑒 𝑒𝑒𝑒
𝑒𝑒𝑒
1 𝑒𝑒𝑒𝑒 𝑒𝑒𝑒
22.4 𝑒 𝑒𝑒𝑒
Gravimetric Factor: the ratio between the grams of a compound and the grams of a single element that
conforms that compound.
Percentage composition: this expresses the mass ratio between a compound's component and the
compound on the basis of parts per hundred
% =
𝑒𝑒
𝑒𝑒𝑒𝑒𝑒𝑒
The percentage composition of a compound may be determined theoretically from the chemical formula
of the compound or from the actual laboratory analysis.
Significant figures are governed by the lab results which were, obviously, measured.
Application of the gravimetric factor
the percentage composition or gravimetric factor of a compound may be used to determine the amount of
a compound's component (ion or element) for an actual amount of compound it may also be used to
determine the amount of compound that may be formed from a given amount of element or ion present
e.g. calculate the mass of cadmium in 25. g of cadmium sulfide
5.4 g CdS multiplied by 112.41 g cadmium, divided by 144.48 g cadmium sulfate comes to 19.8 g of
cadmium
112.41 g cadmium divided by 144.48 g cadmium sulfate, multiplied by 100 is equivalent to 77.803%
Recognize, percentage is based upon 100. In addition, values need units and labels, therefore interpreting
77.803% cadmium in cadmium sulfate would be; 77.803 maa units cadmium any mass unit is okay.
That being the case, that being the case the problem becomes;
25.4 CdS multiplied by 77.803 g cadmium divided by 100.00 g cadmium sulfide, equivalent to 19.8 g
cadmium
NOTE: at times, results may vary slightly between the two methods due to rounding for significant
figures, also, it is permissible to utilize the reciprocals of these ratios as they are still true and correct
Calculating molecular formula:
1) Determine if the empirical formula if not given
2) Calculate the mass of the empirical formula
3) Divide molecular mass by the mass of the empirical formula
4) Round the result to the nearest whole number, multiply it times the subscript for each element of
the empirical formula in order to obtain the subscript for that element in the molecular formula
E.g. An oxide of nitrogen gave the following analysis: 3.04 grams of nitrogen and 6.95 g of oxygen. The
molecular mass of this compound was found by experiment to be 91.0 amu. Determine its molecular
formula
Empirical: 𝑒𝑒2 Molecular: 𝑒2 𝑒4 Empirical mass: 46.01𝑒𝑒𝑒2
Hydrate Formulas: Hydrates are compounds that have molecules of water weakly bonded into their solid
structure in specific ratio. This often occurs as many compounds are formed in solution and isolated
through crystallization. Eating frequently vaporizes and drives off the water of hydration leaving the
anhydrous compound.
It is possible to determine the formula of a hydrate through a similar process. However, the mole ratio f
the water of hydration are included in the calculation.
Eg. Calculate the formula of the hydrate of barium chloride containing 14.7% water of hydration
% BaCl2 = Total % - Water %
= 100.0% - 14.7%
= 85.3% BaCl2
Assuming a 100 g samples
1 𝑒𝑒𝑒𝑒 𝑒𝑒𝑒𝑒2
85.3g BaCl2280.23 𝑒 𝑒𝑒𝑒𝑒2= 0.410 mol BaCl2 / 0.410 mol = 1.00
1 𝑒𝑒𝑒𝑒 𝑒2𝑒
14.7g H20 18.02 𝑒 𝑒2𝑒 = 0.815 mol H2O / 0.410 mol = 1.99
= BaCl2 x 2H2O
Chemistry Notes Unit III Part II
Stoichiometry:
● Chemical reactions and the quantitative relationships between:
○ Reactant and Reactant
○ Reactant and Product
○ Product and Product
Describing a Chemical Reaction:
Indications of a chemical reaction: evolution of heat, light, and/or sound; production of a gas;
formation of a precipitate; color change; difficult to reverse
○ INDICATION not CONFORMATION
○ Precipitate: substances in solid form to be deposited from a solution
Chemical Equations:
● Chemical equations represent chemical change or chemical reactions. They depict the kind of
reactants and products and their relative amounts in a reaction
○ Certain equations provide some quantitative information
○ 4Als + 3O2(g) → 2Al2O3(s)
■ The letters (s), (g), and (l) are the physical states of compounds. Also used are (v), (c)
, (aq)
■ This equation means: 4Al atoms + 3 O2 molecules yield 2 f.u.s of Al2O3 or 4 Al moles +
3 O2 moles yield 2 moles of Al2O3; 4 mol Al @ 27 g/mol, 3 mol O2 @ 32 g/mol, 2 mol
Al2O3 @ 102 g/mol
● 108 g + 96 g = 204 g
● NOT: 4 g Al + 3 g O2 yield 2 g Al2O3
■ The numbers in the front are called stoichiometric coefficients
● Characteristics:
○ The equation represents known facts
○ The equation must contain the correct chemical formulas or chemical symbols for the
reactants and products
○ The law of conservation of mass must be satisfied for formula equations
■ Conservation of atoms
● Reactants enter chemical change; the new substances formed during the chemical changes
○ Reactants left, products, to the right
●
Since a chemical reaction involves the rearrangement of atoms those “reacting atoms” must come into
contact with each other. Te reacting atoms or molecules must collide with one another. If the force of the
collision is sufficient, bonds may be broken and new bonds formed. However, if insufficient energy is
present, the atoms or molecules may merely bounce off each other without a change taking place. This
involves the energy of activation, if it i not met, the reactants can’t get to a point where a reaction can
occur
Chemical reactions are represented by chemical equations. There are several types of chemical equations,
end each one conveys different information
Basic equation structure: Reactants ---> Products
Word equations represent the reaction in terms of names and words instead of formulas and symbols. It
may include some information about the phase or condition of the substances involved in the reaction.
However, it fails to provide information with regard to the quantities or amounts of substances present in
a reaction.
Aluminum hydroxide + phosphoric acid ---> aluminum phosphate + water
A formula or molecular equation uses chemical symbols and chemical formulas to represent the
substances associated with the reaction. This equation may provide information about the phases of the
substances and other conditions within the reaction. Even more importantly, the equation now provides
some information about the relative amounts of reactants and products.
Al(OH)3 + H3PO4(aq) ---> AlPO4(s) + 3 H2O
*State of equilibrium, no NET change is occuring
In a balanced equation, there is conservation of atoms
Skeletal formula equation: unbalanced
Guidelines:
1) Do not start with free elements!
a) Like hydrogen or oxygen, often found in more compounds
2) If a polyatomic ion is found on both sides, treat them as a single unit
3) Think of water as a hydrogen bonded to an OH
4) Some elements appear diatomically, H2, N2, O2, halides, S8, etc etc
5) Check to make sure coefficients represent the simplest ratio
6) ALWAYS write BALANCED AS WRITTEN if all coefficients are ones
Some equations may be balanced using fractions, but the most common approach allows only for integer
coefficients
If polyatomic ions remain intact in a reaction balance them as a group or unit, not individual atoms
NEVEREVEREVEREVEREVER express a coefficient of one
Stoichiometry:
Stoichiometry of reactions - quantitative relationships between reactants and products in a chemical
reaction
Understanding chemical reactions helps is to predict what will be produced by a reaction, it is also
important to know how much will be produced by a reaction, and additionally how much reactant is
required to produce a desired amount of product.
The balanced equation gives the relationship between amounts of reactants and used and amounts of
products likely to be formed, the numerical coefficient tells: how many individual particles are needed in
the reaction on a submicroscopic basis or how many moles are necessary on the macroscopic basis, the
stoichiometric coefficients.
Learning check: For the reaction N2 + 3H2 → 2NH3, how many moles of N2 are used when 2.3 moles of
ammonia are produced?
1 𝑒𝑒𝑒 𝑒2
2.3 mol NH3 x 2 𝑒𝑒𝑒 𝑒𝑒3= 1.2 mol N2
If 0.575 moles of Co2 is produced by the combustion of propane, C3H8, how many moles of oxygen is
consumed; balanced equation C3H8 + 5O2 → 3CO2 + 4H2O
5 𝑒𝑒𝑒 𝑒2
0.575 mol Co2 x
= 0.958 mol O2
3 𝑒𝑒𝑒 𝑒𝑒2
If 454 g of NH4NO3 decomposes, how much N2O and H2O are formed? What is the theoretical mass of each
product.
1) STEP 1: write the balanced chemical equation: NH4NO3 → N2O + 2H2O
2) STEP 2: convert mass reactant (453 g NH4NO3) into moles, use molar mass
1 𝑒𝑒𝑒 𝑒𝑒4𝑒𝑒3
a) 454 g NH4NO3 x
= 5.672 163 9 mol NH4NO3
80.04 𝑒 𝑒𝑒4𝑒𝑒3
3) STEP 3: Convert moles reactant (5.672 163 9 mol) into moles product using the mole ratio from the
stoichiometric coefficients of the balanced chemical equation
1 𝑒𝑒𝑒 𝑒2𝑒
a) 5. 672 163 9 mol NH4NO3 x 1 𝑒𝑒𝑒 𝑒𝑒4𝑒𝑒3= 5.672 163 9 mol N2O
4) STEP 4: Convert moles of product (5.672 163 9 mol N2O) into mass product using molar mass
44.02 𝑒 𝑒2𝑒
a) 5.672 163 9 mol N2O x 1 𝑒𝑒𝑒 𝑒2𝑒 = 249.688 655 6 g N2O = 250. G N2O
Limiting Reactants: If the amounts of 2 or more reactants are given, the reactant used up first determines
the amount of product formed
- It is called the “limiting” reactant: the other reactants are sometimes referred to being in excess
since there will be some left after the limiting reactant is completely used
The limiting reactant is the reactant present in the smallest stoichiometric amount NOT the total amount
- In other words, it's the reactant you will run out of first
The maximum amount of product which can be formed is the theoretical yield
Determining the limiting reagent: There are several approaches to this. One method is to compare the
quantities available to the quantities required.
-
What should be compared? Convert “given” amount of each reactant into moles; if not given as
moles
Apply the mole ratio to reactants in order to compare to the moles available of one of the reactants
to establish the limiting reactant
- Example:
- Ca(OH)2 + 2HCL(aq) → 2H2O(l) + CaCl2(s)
- When 1.00 g of each reactant is combined, what is the theoretical yield of CaCl2
- Limiting reagent? When you are given 2 or more amounts for the reactants, the
first step is determining the limiting reactant
-
1 𝑒𝑒𝑒 𝑒𝑒(𝑒𝑒)2
(limits)
-
2 𝑒𝑒𝑒 𝑒𝑒𝑒
1.00 g Ca(OH)2 * 74.10 𝑒 𝑒𝑒(𝑒𝑒)2= 0.013 495 28 mol (CaOH)2 * 1 𝑒𝑒𝑒 𝑒𝑒(𝑒𝑒)2= 0.0270 mol HCl
1 𝑒𝑒𝑒 𝑒𝑒𝑒
- 1.00 g HCl *
= 0.027 4 mol HCl (excess)
36.46 𝑒 𝑒𝑒𝑒
Use moles or actual amounts of products obtained. Run the stoichiometric calculation for each of
the reactants based upon their “given” amounts. The amount of common product forming that is
the least identifies the limiting reactant and will provide the correct answer.
Chapter 4: Reactions in Aqueous Media
Solutions: Homogeneous mixture in which 2 or more components mix freely
Solvent: Substance which dissolves vs. Solute: Substance being dissolved
Concentration: a solute-to-solvent ratio describing the composition of the mixture
1) % Solution by mass - ratio of the mass of solute to the mass of solution
a) Often expressed as mass solute/100g of solution
i)
Ex: 10.0 g salt and 70.0 g water mixed and the solution is prepared; find the
concentration of solubility by % mass
ii)
𝑒𝑒𝑒𝑒 𝑒𝑒𝑒𝑒𝑒𝑒
10.0 𝑒 𝑒𝑒𝑒𝑒
% concentration = 𝑒𝑒𝑒𝑒 𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒・100 = [10.0 𝑒 𝑒𝑒𝑒𝑒 + 70.0 𝑒 𝑒𝑒𝑒𝑒𝑒]・100 = 12.5% salt
General Concentrations: Concentrations may be expressed very specifically (as with % mass), or very
generally
1) Dilute: Implies “little” solute per amount of solvent
2) Concentrated: Implies “a lot” of solute/ amount of solvent
Solubility: Amount of solute that can dissolve in the specified amount of solvent at a given temperature
(usually g solute/100g solvent or moles solute/L solvent)
Saturated: No more solute can be dissolved at the current temperature in given amount of solvent
Unsaturated: Contains less solute than solubility allows
Supersaturated: Contains more solute than solubility predicts, very unstable solution
1) Most solid solutes are more soluble at higher temperatures, careful cooling of saturated solutions
may result in supersaturated solution, most often forms a precipitate
Ionic Compounds in Water: Water molecules arrange themselves around ionic and dissociate them from
the lattice
The separated ions are “hydrated” and conduct electrical current (act as electrolytes)
*Polyatomic ions remain intact during dissociation
1) Water: Unequal sharing of electrons leads to partial positive charge and negative charges in a
water molecule- charges attract the ions which causes dissociation of the ionic compound in water
a) Molecular compounds: solute particles surrounded, but not dissociated
Electrolyte Properties:
1) Electrolyte: A substance whose aqueous solution forms ions; conducts electricity
a) E.g. ionic compounds
b) Do so by providing ions which are mobile
2) Nonelectrolytes: Substance that does not form ions in an aqueous solution; poor conductor
a) E.g. molecular compounds
Electrical conductivity:
1) Strong electrolyte: Aqueous solution that conducts electricity because solute is 100% dissociated
into ions
a) Exists in a solution completely or almost completely as ions, all soluble ionic compounds
and a few molecular compounds; ex a strong acid
2) Weak electrolyte: Aqueous solution that weakly conducts electricity due to low ionization and low
dissociation
a) Molecular compounds that produce a small number of ions when dissolved in water; ex.
Acetic acid (HC2H3O2) only slightly ionizes when dissolved in water
b) Ionic compounds that are insoluble in aqueous solutions; may produce some dissociated
ions, however, they are few in number
3) Non-electrolyte: An aqueous solution that does not conduct electricity because solute does not
dissociate into ions
a) Molecular compounds that do not ionize entering an aqueous solution- in either case,
conductivity is not enhanced
Ionic equations show dissociate dions: Hydrated ions, with the symbol (aq), are written separately
1) Na2SO4(s)→ 2Na+(aq)+SO4-2(aq)
Symbols in writing chemical equations: (s)- solid, (l)- liquid, (g)- gas, (v)- vapor, (aq)- aqueous media, (c)crystalline solid
↓- precipitation of a gas; ↑- evolution of a gas; Δ- heat applied to reactants
Ionic Equations: Ionic equation represent all soluble ionic salts and ionized molecular compounds as
hydrated ions
Insoluble gases, solids, and liquids such as water are represented in their molecular form or as empirical
formulas, balancing must conserve mass and charge
1) Show only those ions which were changed by the process
a) These are the ions which actually undergo change
2) This is what actually happens in the chemical reaction
a) Eliminates spectator ions: when we compare the reactant to the product, spectator ions
are those which are not changed in any way
i)
Think of fans at a soccer game, they watch not play
Writing ionic equations: Since strong electrolytes exist as dissociated ions in a solution, we can show this
in an equation
1) Identify strong electrolytes
2) Distinguish counting subscripts from characteristic subscripts- counting subscripts become
multipliers
3) Separate the ions into the strong electrolytes
4) Show the states as recorded in the molecular equations
Inorganic Compounds: Frequently classified as either…
Oxides: Combination of any element with oxygen (MgO, CO, N2O3), or, Acids: contain hydrogen for
donation (HCl, HNO3, H2S, H2SO40, Bases: most contain cation and hydroxide ions (KOH, NaOH,
Mg(OH)2), Salts: technically, may be formed from the reaction of an acid and a base- positive ion of base
and negative ion of acid (KCl, Na2SO4, MgS)
Acids: have different definitions pending upon the situation
Arrhenius: An acid increases the concentration of hydrogen ions in the solution. A base increases the
concentration of hydroxide ions in the solution.
● Note- H+ may appear as an increase in H3O+ or hydronium ion.
Brønsted-Lowry: Any substance that donates a proton. A base in any substance which accepts a proton.
Lewis: A lewis acid is an electron-pair acceptor. A Lewis base is an electron pair donor.
Binary Acids/Non-oxyacids:
Naming binary/non-oxyacids: the name of the binary acid consists of two words. The first word has three
components:
1) Part one - the prefix “hydro-”
2) Part two - the root of the nonmetallic element
3) Part three - the suffix “ic”
The second word is always “acid”
Examples: HCl(aq) = hydro chlor ic acid = hydrochloric acid
HBr(aq) = hydro brom ic acid = hydrobromic acid
HF(aq) = hydro fluor ic acid = hydrofluoric acid
Oxyacids: These are more difficult to name because these acids have hydrogen(s), a nonmetal (metal?),
and may have varying numbers of oxygen atoms. How do we name them?
The names of these acids depend upon the oxyanion present: Polyatomic ions ending in “-ate” will
substitute an “-ic” suffix for the “-ate” and the polyatomic ions ending in “-ite” will substitute ab “-ous”
suffix for the “-ite”
If a prefix is attached to the polyatomic anion, it is retained in the acid name.
Examples: HClO4(aq) = hydrogen perchlorate = perchloric acid
HClO3(aq) = hydrogen chlorate = chloric acid
HClO2(aq) = hydrogen chlorite = chlorous acid
HClO(aq) = hydrogen hypochlorite = hypochlorous acid
Properties of acids and bases: The term acid is from acere (latin) meaning sour
Acids have a sour taste, in contrast to bases which have a bitter taste
- E.g. vinegar or lemon juice
- E.g. acorns, kale
- Acids will turn litmus red, while their counterparts the bases turn litmus blue
- Litmus is a vegetable material whose color my be altered depending on the pH of the
environment
- Acids will react with carbonates and bicarbonates to produce carbon dioxide
- 2HCl(aq)+Na2CO3(s)--> CO2(g)+H2O(l)+ 2NaCl(aq)
- Acids and bases are opposites, they will each destroy the properties of the other
- Referred to as neutralization
- Acids and bases will both conduct electrical current
- Therefore, they are placed along with salts in the group referred to as electrolytes (
substances which conduct when placed in an aqueous solution)
- Acids will react with active metals forming salts and hydrogen gas
- Hence acids are frequently considered to be corrosive
- Bases tend to feel slippery because they are caustic
- Caustic: capable of dissolving organic matter
The Arrhenius definition of acid: an acid is a substance that ionizes in a reaction with water to form the
hydronium ion, H3O+; strong acids are 100% ionized when dissolved, whereas weak acids are far less
efficiently ionized
It is common to encounter the hydrogen ion (H+ instead of the hydronium ion, the previous ionization is,
for simplicity also written as HCl(g)→ H+(aq)+Cl-1(aq)
However, H+does not ever exist in aqueous solution, it is always attached to a water molecule as
hydronium.
Non metallic oxides can be acids: nonmetal oxides, or “acidic anhydrides” react with water to form acid
solution: SO2(g) + H2O(l) → H2SO3(aq)
Arrhenius bases: base - substance that produces hydroxide ions in water
Molecular bases undergo an ionization reaction to form the hydroxide ions, and are weak bases
Many nitrogenous compounds are molecular bases
Metal oxides and hydroxides are bases- metal hydroxide solutions dissociate into metal and hydroxide
ions and are strong bases
Soluble metal oxides “basic anhydrides” react with water to form metal hydroxides that are strong bases
Brønsted-Lowry acid must contain at least one ionizable proton!
Weak acids and bases are also weak electrolytes Some acids ionize 100% in water- they are termed “strong
acids” and are also “strong electrolytes”: HCl, HClO4, HNO3, HBr, HI, H2SO4
When the difference between oxygen and hydrogen is two or more the oxyacid is strong
The very soluble metal hydroxides are strong electrolytes and “strong bases”
Group 1 (Family IA) hydroxides and Ca, Ba, and Sr hydroxides - heavy group 2 metal hydroxides
LiOH, NaOH, KOH, RbOH, CsOH
Monoprotic acids: yield 1 proton/ ionization
- Ex. HCl → H+ + Cl- strong electrolyte; strong acid
- HNO3 → H+ + NO3- strong electrolyte; strong acid
- CH3COOH → H++CH3COO- weak electrolyte; weak acid
Diprotic acids: 2 protons - however only the first ionization is complete
- H2SO4 → H+ + HSO4- strong electrolyte; strong acid
- HSO4- → H+ + SO4-2 weak electrolyte; weak acid
Triprotic acids: 3 protons
- H3PO4 → H+ + H2PO4- weak electrolyte; weak acid
- H2PO4- → H+ + HPO42- weak electrolyte; weak acid
Equilibrium: describes a state of no net change
- Two opposing processes are now
occurring at equal rates
- Once established, concentrations do not
change over time, unless impacted by an
external stimulus
- This may be referred to as chemical or
dynamic equilibrium
- Dynamic because opposing
change is continually going
- Often represented by
double arrows in a
chemical equation
- The forward reactions
(left to right forms the ions in a solution and the reverse reaction (right to left)
removes them from the solution
- This equilibrium is present in working with weak acids and weak bases which are
weak electrolytes
- Strong acids and strong bases have a single arrow as 100% of the acid or base is
ionized in solution
- When working with equilibria, the position of equilibrium is often discussed
- If the equilibrium favors the reactants, the position of equilibrium lies to the left
- If the equilibrium favors the products then the position of equilibrium lies to the right
Acid salts: Polyprotic acids can be partially neutralized to form acid salts
- Acidic salts- contains an anion that is capable of furnishing additional hydrogen ions
- The number of hydrogen atoms that can still be neutralized is also indicated in the name of
the salt
A reaction is observed when a precipitate forms a soluble reactant, an acid reacts with a baseneutralization reaction, A weak electrolyte product is formed from strong electrolyte reactants,a gas is
formed from a mixture of reactants
Metathesis: Also called a double displacement reaction or double replacement reaction: General formAB + CD ----> AD CB ; cations change partners but charges on each ion don’t change
Formulas of the products are determined by the charges of the reactant ions
Metathesis reactions occur only if they form a weak electrolyte or non-electrolyte as a product (otherwise,
all ions are spectator ions)
Predicting metathesis reactions: Identify the ions involved
- Do not confuse counting subscripts (those present only to make charges cancel) with those that are
characteristic of a polyatomic ions
- Assign states using solubility rules and knowledge of chemical compounds
Precipitate reactions: insoluble solid that separates from solution
Solubility rules: A general idea as to whether a fair amount of solid will dissolve is achieved using
solubility rules
- All compounds of alkali metals (Group 1A) are soluble
- All salts containing ammonium, acetate, chlorate, perchlorate, and nitrate are soluble
- All chlorides, bromides, and iodides (no FLUORINE), except when combined with Ag+, Pb+2, and
Hg2+2, Cu+
- All sulfates (salts containing SO4-2) except those of heavy alkaline earths
- All metal hydroxides and all metal oxides are insoluble except those of group 1a and of Ca+2, Sr+2,
and Ba+2
- SOLUBLE METALLIC OXIDES are BASIC ANHYDRIDES
- When metal oxides dissolve, they react with water to form hydroxides, the oxide ion does
not exist in water.
- All salts that contain phosphates, carbonates, sulfates, and sulfides, and oxalates are
insoluble except those of Group 1a or combined with ammonium
Acid Base Reactions: Neutralization- metathesis reaction in which acid + metal hydroxide or metal oxide
forms water and salt
NaOH(aq) + HCl(aq) → H2O(l) + NaCl(aq)
Acid base reaction - reaction of weak base and acid transferring a H+
Acid + Base → salt + water + CO 2
Gas forming: metathesis reaction forms one of these products:
- HCN, H2S, H2CO3(aq), H2SO3(aq), NH4OH(aq)
Check for a driving force, formation of weak electrolyte or nonelectrolyte.
Strong electrolytes: strong acids, strong bases, and soluble salts.
Molar Concentrations: In solutions, solutes are dispersed in a larger volume. Molarity expresses the
relationship between the moles of solute and the volume of the solution in liters
- One means of expressing concentration
Molarity (M) = moles solute/VL solution
- Hence, a 6.0 M solution of HCl contains 6.0 moles HCl in a liter of solution
Learning Check: What is the molarity of a solution created by dissolving 10.2 g KNO3 in enough water to
make 450 mL of solution?
𝑒𝑒𝑒𝑒𝑒 𝑒𝑒𝑒𝑒𝑒𝑒
M = 𝑒𝑒 𝑒𝑒𝑒′𝑒
solute
10.2 𝑒 𝑒𝑒𝑒3
1 𝑒𝑒𝑒 𝑒𝑒𝑒3
= 350 𝑒𝑒 𝑒𝑒𝑒′𝑒x 101.10 𝑒 𝑒𝑒𝑒3x
1 𝑒𝑒 𝑒𝑒𝑒′𝑒
=
10−3 𝑒 𝑒𝑒𝑒′𝑒
0.29 M KNO3 - solution ALWAYS identified by its
Making a Molar Solution:
→ how to prepare a molar solution
1) Based upon the molar concentration and the volume of solution desired, determine the moles of
solute to be dissolved
2) Convert moles of solute needed into mass in order to obtain a measurable quantity
3) Obtain the desired mass of solute and place into a volumetric container
4) Add solvent (frequently water) till 3/4th of the container is filled
5) Agitate to dissolve the solute
6) Add additional solvent to obtain the desired volume
7) Agitate again
Dilution: Adding solvent to a solution creates a less concentrated solution. Moles of solute present do not
change, hence;
MstockVstock=MnewVnew
- M = molar concentration
- V = volume
Using volumetric glassware ensures that the volumes are known precisely
Alternate equation: Mi x Vi = Mf x Vf -- i - initial, f - final
Adding solvent does not change how many moles of solute are present, the total volume of solution does
change, the concentration of the solution is decreased while the actual amount of solute is unchanged
A volumetric pipette is used to transfer stock solution - a volumetric flask is used to receive final solution
Solution Stoichiometry: A balanced chemical equation is needed to start any stoichiometry problem
In addition, if we are given starting quantities of more than one reactant, one must determine the limiting
reagent
The difference arises in how we calculate moles of reacting substance
When it is a volume of a molar concentration we know moles solute = Vl sol’n x M
1) Step 1 - balanced chemical equation
2) Step 2 - convert given to moles
3) Step 3 - apply stoichiometric mole ratio
4) Step 4 - convert to desired unit (volume) using the molarity
Ion Concentrations:
- The chemical formula for a strong electrolytes relates the moles of ions that will be released on
dissociation to the chemical formula.
- Thus the formula can be used to relate the ion concentration to the solution concentration
Titration: Titration is the controlled addition of one reaction (titrant) to a known quantity of another
(titrate) until the reaction is complete
- Often an indicator is used to signal the reaction completion
- Endpoint: the volume of titrant required to complete the reaction
Solving titration problems:
1) Write the balanced equation
2) Calculate the moles of the known component
a) Moles = Vl sol’n x M
b) = mass / molar mass
c) = particles / 6.022 x 1023
3) Use stoichiometry to determine moles of the unknown
4) Convert moles to desired quantity