PHYSICAL SCIENCE
SLT STUDY GUIDE
Chemistry and Physics
2014-2015
Mass vs. Matter
• Matter is something that occupies spacetime
and has length, width and height (and time)
• Mass, on the other hand, is HOW MUCH
MATTER an object has at any one time
– It is a measure of inertia
– Measured in grams, milligrams or kilograms, etc.
– The mass of an object does not change
– Mass and Matter follow the Conservation Laws –
they can’t be created nor destroyed – only
transformed
• Mass remains the same regardless of the
phase of the matter
• However – the volume may change due to the
substances density in that state of matter
1 kg
1 kg
1 kg
Small Volume
High Density
Medium Volume
Medium Density
High Volume
Low Density
• NOTE: Mass is NOT the same as Weight
• Weight is mass with the effect of gravity
applied to it!
– Wt. = mass x gravity (9.81 m/s2 on Earth)
• Mass does not change – gravity does change
from place to place –
so weight will change
Volume and Density
• Volume is how much space a certain mass
occupies – typically calculated by multiplying
width x height x length of the container
– Measured in cm3, L, ml, etc.
• Density is how much mass is compacted into a
space
– In other words, how closely packed are the atoms
of the matter
– D = mass / volume
– Answers will be in g/ml, g/L, etc.
Types/Phases of Matter
• The type or phase of matter is dependent
upon the KE in the particle’s of that matter
• All particles (atoms or molecules) have KE
• They move by rotating, vibrating or translating
(going in a straight line)
• So, the amount of KE determines the phase of
the matter
– Particle KE is affected by things like temperature
and pressure
1. PLASMA
• On the scale of KE, this phase/type of matter
has the highest amount
• In fact, it has so much, the electrons get
ripped off the atom’s nucleus – leaving plasma
• Example is lightning or the surface of the sun
2. Gas
• Gases are next on the KE scale
• A gas is defined as a type of matter that fills its
container, so its volume equals whatever it’s in
• Gases expand to fill their container in a
process called diffusion
• Since the atoms of a gas have a lot of space
between them, gases can be condensed or
compressed
– In fact – applying pressure and lowering the temp
of a gas will transform it into a liquid
• Gas particles – due to the space and KE –
collide a lot
– They collide with each other and the sides of the
container they are in
– The collisions with the container we measure and
call pressure
– The rate of collision is
the gas pressure in the
container!
3. Liquid
• Like gases, they fill the container they are in –
so their volume is measured by their container
– A liquid’s mass and volume can be measured
separately
• However – there is less space between the
particles and they are less compressible
Key Terms with Liquids
• Viscosity – this is the rate
or ability of a liquid to flow
• Surface Tension – the top of a liquid has a
“skin” on top due to the attraction of the top
layer of particles by the ones below it
– This is what allows a drop to form a circular shape
• Vapor Pressure - due to evaporation, there
will be some particles of a liquid in the
atmosphere above it, and this changes the
pressure above that liquid
• Capillary Action – this is the ability of a liquid
to flow uphill!
– It is achieved through electromagnetic forces!
4. Solids
• Much lower particle KE and this is limited
mainly to vibrations and rotations
• Solids have a set volume, mass, shape and
density
• Just like the other types of matter – solids are
affected by temperature and pressure
5. Bose-Einstein Condensate
• This form of matter can only be achieved at
absolute zero!
– That’s -465oF, -273oC, or OoK!!!!!
• This means that there is NO KE in or between the
matter’s particles
• Because of this – light sent through BEC will slow
down to 7, that’s right, 7 mph!!!!!
• Temp lowered by using lasers
Phase Diagrams and Concepts
• Phase diagram – a graph to show the phases of a
substance and at what temperature and pressure it
goes into each phase
• Phase Changes:
From 
Solid
Liquid
Liquid
Gas/Vapor
Solid
Gas
To
Liquid
Solid
Gas/Vapor
Liquid
Gas
Solid
Term
Liquification / Melting
Solidification / Freezing
Evaporation
Condensation
Sublimation
Deposition /
Crystallization
Classifying Matter
Baryonic Matter
(Observable Matter – 20% Universe)
Substances
Mixtures
Element
Compound
Heterogeneous Homogeneous
-On Periodic
Table
-Pure Form
-Single group
of Atoms
-2+ elements
-Molecules
-Proportional
make up
-Separable by
filtering, distilling
or crystalizing
-Suspension
-Emulsion
-Includes:
Colloids and
Solutions
-Non-separable
Matter Concepts:
• Element
– The pure form of a substance
– Found on the periodic table
– Can be in any of the phases
mentioned!
• Compound
– A combination of 2 or more elements
– Combination is in a ratio or proportional
– Basically – a molecule
• Mixtures
– 2 or more elements or compounds mixed together
– Heterogeneous – a mixture that can be easily separated
through filtration, distillation or crystallization
– Suspension – a temporary mixture where the particles
separate out on their own (like dirt in water)
– Emulsion – a suspension of two liquids – will separate
(mayonnaise)
– Homogeneous Mixture – non-separable, remain
mixed completely and equally
– Includes Colloids – a solution with large particles
dissolved in it, like milk or paint
– Includes Solutions – a solution is made up of the
solute (what is dissolved) and the solvent (what
the solute is dissolved in)
• Example: Salt Water  Salt is the Solute and water is
the solvent
Chemical and Physical Properties
Physical
Does not change with phase
Density, Mass
Boiling/Freezing Points
Color
Hardness
Odor, Taste
Texture
Malliability
Ductile
Chemical
May change in a reaction
Atomic Structure
Flammability
Oxidation
Reactivity
Bonding
Radioactivity
ATOMIC MODELS
Protons (+ charge) and Neutrons (0 charge) make up the nucleus;
Electrons (- charge) around nucleus and held by electromagnetic
force. p+ and n made up of quarks.
Plum Pudding Model
Electron Cloud Model
Scientists
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Democritus
Aristotle
J. Dalton
W. Crookes
JJ Thompson
R. Millikan
E. Rutherford
J. Chadwick
First to believe matter made of atoms (atomos)
Matter made of four elements (fire, air, earth and water)
-Used by Church for nearly 1500 yrs.
-Experiment to gain knowledge
-Gases made of corpuscles (atoms) w/ spaces
Atomic Theory
-All elements made of atoms
-All elements of same kind have same atoms
-Reactions are changing, etc. of atoms
-Conservation of Matter
Used CRT to find electrons
Determined charge on electron to be negative(magnets on CRT)
-Plum Pudding Model of Atom
Mass and charge amount of electron
-Oil Drop Experiment
Gold Foil Experiment
-Used alpha radiation (a He nucleus) to find protons
-Beta and gamma radiation
Found neutron (no charge)
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N. Bohr
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Standard Model of Atom
Einstein
E. Schrodinger
W. Heisenberg
P. Dirac
L. De Broglie
Four Forces
Strong
Weak
EMF
Gravity
Quarks
Atoms give off certain colors/wavelengths of light
-Planetary Model of Atom
Brownian Motion – atoms shown to exist by hitting pollen
Photoelectric Effect – light is wave or photon
Light (photon) is a packet of energy w/ no mass
Electrons at speed of light
Electron Cloud Model
Uncertainty Principle – never certain of election place
Light used to see electrons makes them move
Anti-matter
Electron a wave and particle too
Holds nucleus together with gluons
Hold each neutron and proton together (quarks) w/ bosons
Holds nucleus to electrons - photons
No one knows – maybe a graviton
Includes subatomic particles (260+)
Developed from Quantum Theory
All matter made of Fermion particles
Baryonic Matter made of quarks (Hadrons -3; Mesons – 2)
Non-Baryonic includes leptons (electrons)
Make up Baryons
3 in proton and 3 in neutron
2 in Mesons
6 Types – up, down, top, bottom, charm, strange
Four Universal Forces
Standard Model
 Pauli Exclusion Principle – No two e- in the same
atom can have the same quantum numbers
In each suborbital, there will be two e-,
and each will have opposite spins
 Aufbau Principle – Atoms will fill up their orbitals
starting with the lowest energies to highest
 Hund’s Rule – Each suborbital will have a
maximum of 2 e-; with opposite spins. Each
suborbital in an energy level must have one ebefore any receive a second one
Valency
• Valence Electrons – the electrons in the
outermost shell or energy level
– Octet Rule – atoms seek to fill to completion the
outermost energy level by gaining or giving away
electrons
– Valence electrons will determine bonds and
reactions
– Valency provides grouping of elements on Periodic
Table (Groups I through VIII)
Electron Configuration
• We need to be able to show how an atom’s
electrons are configured (distributed) in order
to explain chemical bonds and reactions
• There are several ways to do this, but first:
– You must know and understand the element’s
valence energy level (the period)
– You must know the # valence e- (Group I to VIII or
if a Transitional Metal)
– And you must know the orbital grouping: s, p, d,
or f
Aufbau Notation
• This gives the valence energy level, the valence
orbital shape and the # of valence electrons in
that level
Example:
Energy Level
2p5
Orbital
# e-
Aufbau Notation:
• Long Format:
– Na w/ 11 electrons would be: 1s22s22p63s1
– Short Format only gives the LAST Energy Level and
the number of electrons in it
• Na  3s1
– Without this e-, Na would isoelectric w/ Ne (have the same econfiguration)
• Cl -17 would be 1s22s22p63s23p5 in Long Format
• Cl – 17 would be 3s23p5 in Short Format and its Lewis
Dot Structure would have 7 dots around the elemental
symbol
• Often, you will see another shorthand way of
writing an element’s electron configuration
– This is called the Noble Gas Configuration
– Noble Gases are in Column VIII
– For example, using Na, Na’s configuration would
be: [Ne]3s1
– The [ ] means isoelectric (the starting point)
– Li would be [He]2s1
- La  [Xe]6s25d1
– Mn would be [Ar]4s23d5 - Ce  [Xe]6s25d14f1
– Zn would be [Ar]4s23d10 - Pr  [Xe]6s24f3
Lewis Dot Diagrams
• Dot diagrams demonstrate the type of
covalent bonds an element may make under
certain conditions
• The element’s symbol is surrounded by up to
8 dots, each dot representing a valence e• Maximum of 2 dots per side (2 dots x 4 sides
= 8 total, an octet valence level)
• It can be used for single atoms or to show
molecules and their bonds and shapes
• Determine dots from either doing the e- mapping
OR learn to read the periodic table!
• Often, for transitional metals – need to do
mapping (problem if ion)
• Remember Hund’s Rule
• Watch for + and – ions and either add or subtract
the correct number of electrons
• Each side of elemental symbol represents one of
the last 4 suborbitals (□) of the SAME Energy
Level!
• Do not have to have 8 e- dots, can have openings
– This is where bonds to other atoms can occur
– Note: RN columns all have same dot drawing!
Lewis Dot Diagrams
Nucleosynthesis
• Due to the super high temperatures and density, the e-, e+, p+,
and n0 began to collide at the start of the universe
– They got close enough for the Strong Force to take over
allowing nuclei to form
• The first element was the nucleus of an isotope of Hydrogen,
Deuterium (2H) – a p+, n0
• This was joined by a p+ and formed 3He
• Then, with another n0, 4He formed
(the stable form)
• With the capture of free e-, a stable atom was made
• Process called alpha combination
Nucleosynthesis in sum:
When dealing with Gases, there
are four major variables utilized:
Volume
Pressure
Temperature
Amount of Gas (Moles)
Gas Equation Summary
• Boyle’s Law
 If P goes up, V goes down (inverse)
• Charles’s Law
 If V up, T up
• Gay-Lussac’s Law
P1V1 = P2V2
V1 = V2
T1 T2
P1 = P2
T1 T2
 If T up, P goes up
• Combined Gas Law
P1V1 = P2V2
T1
T2
Dmitri Mendeleev (1834-1907) is given usually given
credit for developing the first periodic table based on
atomic weight and properties – it allowed him to
predict new elements.
His predictions were proven true with the discovery
of gallium, scandium and germanium
However, it was Henry Moseley (1887-1915)
who set it up using the atomic number (Z)
He used x-rays to note the frequency of the
element as well as the number of
p+
The chart became arranged
according to atomic number
rather than atomic weight
ISOTOPES
• All atoms of an element have the SAME
number of protons (p+)
• The p+ number is the atomic number (Z)
– This is a constant – it stays the same for that
element’s atoms
– For example: All Sodium (Na) atoms have 11 p+
– If an atom loses a proton, it becomes a different
element
• If Na loses 1 p+, then it has become Neon (Ne)
Z = atomic number = p+
• The number of protons identifies the atom
and which element it is
• In a stable atom:
– # p+ = # n0 = # e– Thus, Na in its stable form has 11 p+; 11 n0; and 11
e– If it has an unequal number of p+ and n0, then it is
called an ISOTOPE
The Carbon Isotope
IONS
• Ions are when an atom has an unequal
number of p+ and e• Metals form (+) ions and nonmetals (-) ions
• Remember – a stable atom has a neutral
overall charge due its equal number of p+ and
e• When an atom loses or gains an e-, its charge
changes accordingly
– Loss of e- means a + charge; gaining an e- means
a – charge for the atom
Losing or Gaining e- . . . . .
• If an atom loses an e-, then it has more p+ than
e- and it will have an overall positive charge
• Different elements’ atoms can lose 1, 2, 3, or
even 4 electrons depending on various factors
• If an atom has LOST e-, then it is called a
CATION or a positive ion
– A Cation would be written as Al+ (the one being
understood) or Al+3
• Atoms can also gain electrons
• If an atom gains electrons (from 1 up to 4), then it
will have more e- than p+ and will end up having
an overall negative charge
• A negatively charged ion is called an ANION
• A positively charged ion is called a CATION
• The NOBLE GASES will not form ions and thus will
not bond
• The Transition Metals can form various numbers
of positive ions – got to learn these!
• The losing or gaining of electrons determines
what type of bonds the atoms will form, and
which atoms will bond to others
Using the Periodic Table
• Elements in the Main Groups (A), form fairly
consistent ions
• Group I will form +1 ions; Group II form up to +2;
Group III form up to +3 ions
• Group IV will form either up to -4 or +4 ions
• Group V will form up to -3 ions; Group VI up to -2;
Group VII form -1; and Group VIII will not form ions
at all
• Those elements in the B group (transition metals)
vary in their + charges meaning they can form
different ions – look them up!
Group Names - Periodic Table
Know These!
Ions and Isotopes in Review
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Stable atom: #p+ = #n0 = #eAtomic Mass: #n0 = # p+
If charge is 0, then #p+ = #eIf charge is positive, then #p+ > #e- Cation
If charge is negative, then #p+ < #e- Anion
Mass measured in AMU (Atomic Mass Units)
based on the C-12 atom
Examples:
• Li-1 has gained an electron, meaning there is
one more negative charge than positive ones
– It has 3 p+ and 4 e-
• Li+1 has lost an electron, meaning there is one
more positive charge than negative ones
– It has 3 p+ and 2 e• REMEMBER: The # of p+ DOES NOT CHANGE
• Only the number of n0 (isotope) and e- (ion) change
• Cf-3 has an atomic number of 98
– This means it has 98 p+
– Its atomic mass is 216
– It has 118 n0, (216 – 98), making it an ion and an
isotope!
– Since it has a -3 charge, the number of e- will be
101; (98 + 3)
– Zn+1 has 30 p+ and n0; but due to the +1 charge, it
has only 29 e-
On the Periodic Table:
The top number is Z, the Atomic Number or
number of p+
The Element’s Symbol
The element average atomic weight set by isotopes
and abundances
Counting Atoms in a Molecule
In the example, NH3, the subscript 3 only applies to the hydrogen.
– Therefore: there is 1 N and 3 H in ammonia
In the example, 3Ca3(PO4)2, the number of atoms changes due to the
Coefficient is always in front of the whole molecule!!
-The subscript 2, multiplies the P (2) and O (4 x 2 = 8) since it is
outside the parenthesis
-The subscript 3 only goes with the Ca
-The coefficient 3 is multiplied to the Ca, P and O after you do the
subscripts
-Therefore, this molecule has (3 x 3) Ca + (3 x 2) P + (3 x 4 x 2) 0
which equals 39 atoms
3Ca3(PO4)2
Ionic Bonds
• These are the bonds between a metal and a
nonmetal
• The metal Ion is positively charged and called
a cation
• The nonmetal Ion is negatively charged and
called an anion
• The bonded molecule should be neutrally
charged when finished
Knowing where the metals and nonmetals are on the
table will make your life easier
Covalent Compounds
• These can be monatomic or polyatomic
compounds
• It is a bond between two nonmetals
• They share a pair of electrons
• They can be subgrouped into polar or
nonpolar
• If a binary compound (2 atoms) – use the
same naming rules as in Ionic Compounds
Naming Covalent Compounds
Process:
1. Prefix Indicating # + full name of first
nonmetal
2. Prefix Indicating # + root name of second
nonmetal + suffix “ide”
3. Watch for polyatomics and use their proper
names
• If it has more than two atoms – need to use
the prefixes
Number Prefix
1
Mono
2
Di
3
Tri
4
Tetra
5
Penta
6
Hexa
Number
7
8
9
10
11
12
Prefix
Hepta
Octa
Nona
Deca
Undeca
Dodeca
For Example:
NOTE THESE ARE ALL NONMETALS WITH NONMETALS!
• P4S10 becomes Tetraphosphorous Decasulfide
• P2O5 becomes Diphosphorous Pentaoxide
• SF6 becomes Sulfur Hexafluoride
• N2O3 is Dinitrogen Trioxide
• CO is Carbon Monoxide
• SO2 is Sulfur Dioxide
• SiBr4 becomes Silicon Tetrabromide
• Water is really Dihydrogen Monoxide!
Naming Ionic Compounds is really simple:
1. Name the cation (metal) using its proper name; if it is a
polyatomic (NH4+1 ammonium), do the same
2. Then, using the stem of the anion (nonmetal
mono/polyatomic), simply add the suffix “ide” to it
3. If a transition metal with different possible ions, a roman
numeral will tell you which one it is – and it changes the
molecular formula! [Iron (III) Sulfide]
Examples:
Iron (II) Sulfide = Fe+2 and S-2 combined
Zinc + Chlorine = Zinc Chloride
Iron + Oxygen = Iron Oxide
Lithium + Cyanide = Lithium Cyanide
Ammonium + Fluorine = Ammonium Fluoride
Cobalt + Phosphorous = Cobalt Phosphide
Balancing Compounds
In an Ionic Compound – balance the molecule
using the criss-cross rule. Switch oxidation
numbers, making them into subscripts and
DROP charges.
Mg
+2
+
-1
Cl
Mg Cl2
The one is understood.
This applies even if using a polyatomic ion
NH4+ +
O-2
(NH4)2O
The parentheses are used to keep
the polyatomic together; the 1 is understood
Pb+4 +
CO3-2
Pb2 (CO3)4
Pb(CO3)2
and this can be simplified by
reducing the subscripts to
Chemical Equations
• The chemical equation is the rxn formula
• Reactants  Products
– Each component will have a phase indicator:
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(g) meaning it is in its gaseous phase (not just gassy)
(l) meaning it is in its liquid phase
(s) in its solid phase
And (aq) meaning the substance is in a solution of
water, aq meaning aqueous
• Must remember which elements are normally
diatomic (N2, O2, F2, Cl2, Br2, I2, and H2)
– AND S8 and P4 if by selves!!!!!
• All molecules in an equation must be balanced
first!!
– Remember the criss-cross rule!!
• You may not adjust any subscripts from the
original formula
• You may add and adjust, as you will see, the
coefficients in front of each item in the
equation
Example:
• H2 + O2  H2O
– This is the skeleton equation
– According to the Law of Conservation of Matter,
both sides of the arrow must have the SAME
number of atoms for each and every element –
NO EXCEPTIONS
– The  can be treated like an = sign
– In reality, it indicates that some sort of process
occurred to cause the reaction
– So. . . . .
• To balance this simple equation:
– We ARE NOT ALLOWED TO CHANGE SUBSCRIPTS
– We CAN ADJUST COEFFICIENTS ONLY
– The subscripts are the numbers after and below
each element’s symbol
– The coefficients are number in front of a unit
(atoms or molecules) and tell how many units
there are
– The coefficients are multiplied out to each and
every unit’s atom they are in front of
– So. . . . .
H 2 + O 2  H2O
• There are 2 H and 2 O on the reactant side of
the equation (the left side)
• There are 2 H and only 1 O on the product
side (the right side)
• Each side must balance
• You may add, adjust, finagle, cram, etc. any
coefficient in front of any and/or all units to
get the equation to balance
• Therefore: 2 H2 + O2  2 H2O
2H2 + O2  2H2O
• Now this is balanced!
• It means it takes 2 hydrogen molecules and
one oxygen molecule to form 2 water
molecules
Balancing Equations Steps:
• First identify all the reactants and products in
the equations
• Remember – subscripts indicate how many of
each element’s atoms are present – with 1
being understood
• Remember to multiply out all subscripts that
are outside a unit in parentheses!
• YOU CAN’T CHANGE SUBSCRIPTS
• COEFFICIENTS HAVE TO GO IN FRONT OF A
UNIT
• Let’s take the unbalanced equation of:
KClO3  KCl + O2
• List the elements and how many for both
sides of the arrow
K 1  K 1
Cl 1
Cl 1
O 3
O 2
• Obviously, everything is fine except for oxygen
– This is where we have to adjust
• We can only use coefficients
– So we try to multiply each Oxygen by a number to
get them to equal out
– These multipliers become coefficients
K 1
 K 1
Cl 1
Cl 1
O 3x2=6
O 2x3=6
• So the new equation is:
2 KClO3  KCl + 3 O2
• This changes the number of K and Cl now
• You have to readjust again. . . . . .
2 KClO3  KCl + 3 O2
• Now we have:
K 2  K 1
Cl 2
Cl 1
O 6
O 6
• Multiply the product KCl by a coefficient of 2 and it
balances
• Let’s check:
2 KClO3  2KCl + 3O2
K 2  K 2
Cl 2
Cl 2
O 6
O 6
• It’s Balanced! Finally.
Reaction Types
• SYNTHESIS (or Direct Combination or
Composition) REACTIONS
– 2 + reactants join together to form a single
product
– Resulting compound is based on common
oxidation numbers of the reactant elements
– There is typically an electron transfer from the
element with the lower EN to the one with the
higher EN
– So: A + B  AB or AB + C  ABC
– If two nonmetals involved – a covalent bond
formed
– If two metals – a metallic bond
– If metal with a nonmetal – ionic bond
• DECOMPOSITION REACTIONS
– Compounds break down into components
• AB  A + B
or
ABC  AB + C
• Examples. . .
CaCO3  CaO + CO2
2 KClO3  2 KCl + 3 O2
H2CO3  H2O + CO2
Ca(OH)2  CaO + H2O
2 NaCl  2 Na + Cl2
• REPLACEMENT REACTIONS (2 types)
Single Replacement (Displacement Rxn)
– Key Rule: Metals Replace Metals
• A + BC  AC + B
– If Nonmetal – a transfer of e- from more reactive
to lesser one
– Halogens Replace Halogens also
– Metals replace H in H2O  Metal OH- + H2 (g)
– Metals replace H in Acids  salt + H2(g)
• Al + H2SO4  AlSO4 + H2(g)
• 2 Sc(s) + 6 HCl (aq)  2 ScCl2(aq) + 3 H2(g)
• DOUBLE REPLACEMENT
Example:
– FeCl3 + 3 NaOH  3 NaCl + Fe(OH)3
OH goes with FE
Cl goes with Na
• Cations exchange anions with each other
– No change in oxidation numbers
– Better know your ions and polyatomics
– Remember the criss-cross rule and balance each
compound after exchanging anions!
– So: AB + CD  AC + BD
• COMBUSTION
– An exothermic rxn (gives off energy)
– Usually find CO2 and H2O in products
– O2 usually found in reactants
• CH4(g) + 2 O2  CO2(g) + 2 H2O(g) + heat
• 2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(g)
• ACID/BASE REACTIONS
An acid + base  salt + H2O
– Acids lose a H+ ion and the bases lose OH- ion
• These make up one of the products, water
– Process is called neutralization
– The produced salt does not have to be NaCl and
can be any ionic compound
– Measure acid with pH scale (1 strong, 7 neutral
and 14 is a base)
– Measure base with pOH scale
• Exothermic Reaction – Gives off energy during
reaction
• Endothermic Reaction – Absorbs heat/energy
during reaction
Solutions (solns)
• Definition:
– It is a homogeneous mixture
– It is made up of a solute (what is being dissolved)
and a solvent (what the solute is dissolved in)
• Solute + Solvent = Solution
– It can be a gas, liquid, or a solid or even a
combination of these phases
– It does not separate into its parts on its own
Solubility
• If a solute dissolves in a solvent, it is soluble
– Solubility affected by Temperature, Pressure,
Agitation, etc.
• The process in which the solvent particles
surround the solute and “cause” it to dissolve
is called solvation
– If this occurs in water – called hydration
– Rule of Solvation – “Like Dissolves Like”
• Rule of Solvation
– “Like Dissolves Like”
• Refers to polar solutes dissolving in polar solvent, but
not in nonpolar solvents
• This is due to the charges found on the molecules and
opposite charges attracting each other
• If it does mix – said to be miscible, and if doesn’t –
immiscible
– Immiscible mixtures are not solutions since they will filter or
separate on their own
– Like oil and water
• If the solute is an ionic compound and is
dissolving completely, it is breaking down into
its cations and anions
– This is called ionization
– Dependent upon the dissociation energy (Do), the
energy needed to break apart the ionic compound
– This is a part of the solvation process
Saturation
• Saturation refers to the level of solute in the
solvent
• If the solvent can not dissolve any more
solute, then the two are in equilibrium and
the solution is saturated
• If the solute “falls” out of the solvent –
oversaturation
• If not enough solute - undersaturated
Solution Types
• Suspension - the solute will filter out of the
solvent on its own (dirt in water)
• Thixotrope - the solute and solvent adt as a
solid until agitated, then it acts as a liquid
• Emulsion - a liquid with a solid solute
dissolved in it
• Aerosol - a colloid using a gas as a propellant
• Colloid - a heterogeneous mixture, the solute
does not settle from solvent
Colligative Properties of Solns
• Properties of solns that depend on the
number of solute particles in the solvent
1. Vapor Pressure – affecting evaporation rates.
This is amount of solvent found in gas phase
above the solution. Lowers vapor pressure.
2. Boiling Point – more solute, higher the Bp
3. Freezing Point – more solute, lower Fp
4. Osmotic Pressure – affects the passing through
of semi-permeable membranes
Solutions Summary
• Soln = Solute (what is being dissolve) in a solvent (what it is being
dissolved in)
• Solute Effects the Vapor Pressure (evaporation rate); the boil and freeze
points and the osmotic pressure of the solution
• Saturation vs. Dilution
– Precipitate
• Not easily filtered (based on particle size)
– If mix – called miscible
– “Like Dissolves Like” refers to polar and nonpolar substance mixing
– Other Common Solns:
•
•
•
•
Suspension – larger particle size means will filter out
Emulsion – a solid solute in a liquid solvent
Thixotrope – acts as a solid or liquid until agitated
Colloid – hetereogeneous mixture where solute will not entirely filter out
Acid (pH) and Base (pOH)
• Strong Acid (1)
• Neutral (7)
• Strong Base (14)
A/B General Info
• Bases
–
–
–
–
–
Also called alkaline solutions
Have OH in them
React with acids in neutralization rxn to makes salts and water
Turn pink litmus paper blue
Taste bitter and feel soapy
• Acids
–
–
–
–
–
Usually have H in them with (aq) after their name
Ionize in water
React with bases to make salts and water
Turn blue litmus paper pink
Taste sour
• A/B Rxn
Acid (with H) +
H w/ nonmetal
Base (with OH)
Metal with OH
 Salt
+ H OH
 Metal (from B)Nonmetal(from A) +
H2O
Other Notes:
– pH + pOH will equal 14
• pH and pOH are INVERSE to each other
• pH measure ionization of H+; pOH measure ionization of OH• Dissociation of ions determines strength of A or B
– As acid gets stronger, base weaker – an inverse relationship
– There are 7 main Bases (most have OH in them)
END OF CHEMISTRY REVIEW SECTION!!!
Start of Physics
Review Section
Speed, Velocity and Acceleration
• The speed of an object is the distance the object
travels per unit of time. Speed is a rate which tells
you the amount of something that occurs or
changes in one unit of time.
• Speed=distance over time
• Speed can be divided into two subtitles constant
speed & average speed. Constant speed is the
speed that does not change. Average speed is the
total distance divided by time.
• Velocity is a speed in a given direction
Velocity
• V1 represents the initial or starting velocity
– If the object starts from a rest, V1 will = 0
• V2 represents the final velocity of an object
– If the object ends with a stop, then V2 right at the
end will be a zero, but not just a millisecond
before that!
–V = d/t
– And this means d = vt; and t = d / v
Acceleration
• The acceleration of an object as produced by a
net force is directly proportional to the
magnitude of the net force, in the same
direction as the net force, and inversely
proportional to the mass of the object.
• Acceleration (a) = ΔV / Δt -or• Acceleration = force over mass
Newton's 1st Law of Motion
• An object at rest tends to stay at rest and an object in motion
tends to stay in motion with the same speed and in the same
direction unless acted upon by an unbalanced force.
• Sometimes referred to as the “Law of Inertia."
– Inertia is the state of rest or resisting a force that may cause motion or
a change in velocity
• Frame of Reference – how the observer sees the change in
velocity
• Frame of Reference (Point of View) can be stationary or
moving depending on the observer
• Example: When a car stops suddenly, all the loose objects will
continue forward until they hit something that stops them
(have you ever had coffee do this at a stoplight?)
Newton's 2nd Law of Motion
• The second law states that the acceleration of an
object is dependent upon two variables - the net
force acting upon the object and the mass of the
object.
• It explains the relation of force, mass & acceleration.
• Force=mass x acceleration (F = ma)
• Weight is also a force = m x g
• The net force on an object is equal to the product of
its acceleration and its mass.
Force
• Force is measured in the SI unit called a
Newton (N)
– 1 N = 1 kg x 1 m / s2
1 N = .225 lbs
1 lb. = 4.448 N
• Forces usually are in equilibrium (balanced)
• Weight is a Force (wt = m g)
Force Continued. . .
• By definition it is a push or pull
• It can be divided into two subsets: unbalanced
and balanced
• Unbalanced force can cause an object to start
or stop moving; or change its acceleration,
velocity or direction
• A balanced force is equal forces on an object
that will not change the object’s motion
Acceleration - Due to Gravity
• agrav or just plain g, has a value of 9.80665
m/s2
– We’ll round this off to 9.81 m/s2
• Use 10 for guesstimating!
– Believe it or not – agrav at the equator is
9.7804 m/s2 and at the poles it is 9.8321
m/s2
Free Fall Acceleration
• If v1 (initial velocity is zero or the object is at
rest then falls):
– V2 = gt
– V2 = √2gh
– H = ½ gt2
– H = v2 t
2
If v1 does not equal zero. . . the object
is thrown down or is shot downwards
• V2 = v1 + g t
• V22 = v12 + 2 g h
• H = v1t + ½ g t2
• H = v2 + v 1 t
2
Momentum (ρ)
• Momentum is the product of an object’s mass
and velocity
• It is directly proportional to mass and velocity
• It’s the tendency for an object to keep in
motion
–p = mv
– F t = m v; where F t is the impulse or change in
momentum
Newton’s 3rd Law
• Basically – the law means that for every action
there is an equal and opposite reaction
• A rocket launch – the Fthrust downwards
(action) forces the rocket upwards (reaction)
against the Fgravity
•
Remember:
Action  Equal/Opposite Reaction
Vectors
•
Properties of Vectors
–
–
–
–
–
Vectors can be rearranged into a diagram
Size and Direction can not be changed
Use the tip to tail method of rearranging vector arrows – addition
To subtract vectors – add one to its opposite
Gives final displacement
– Example: One Dimension
A
(5 m) + B
(6 m) = R (Resultant) of 11
m to the right (same direction = addition)
A
(4m) + B
(-3m) = R of 1 m to
the left (opposite directions = subtraction)
Some Key Concepts
• Mass – the amount of matter something has
• Weight – mass affected by the force of gravity
(m x g) – this a Force!
• Density – how much mass per volume
(d = m / volume)
Can be determined through math or
through the displacement of fluid method
*(Remember Archimedes and the crown)
Work, Power, Energy
Work
• Work is a force applied to an object that causes displacement
• W = FΔd
– Measured in Newton-meters (Nm) or Joules (J)
– Kg m2 / s2 is also called a Joule
Power
• Power is the rate at which work is being done
• Measured in Joules per Second or Watts (W) and 1 J/s = 1 W
• Power (P) = Work / Time
– P = W/t = Fd/t
Energy
• Potential Energy
– The stored energy of position, inertia, or ability to do work
– PE = m g h
• Kinetic Energy
– Energy of motion
– KE = .5 m v2
– KE = F d
KE / PE Example
• As KE increases, PE decreases and versa vice
• As it moves upwards, height increases and PE increases; as it moves
downwards, velocity increases and KE increases
• The pendulum at the two highest points have high PE, but no KE until it
starts to move towards the center again.
• Then, the PE decreases until the bob hits the bottom and KE is at its
highest
•No medium needed
•All are transverse waves
•Have an electrical and magnetic field at right angles to each other
•Longest Wavelengths  Shortest λ
•Lowest f  Highest f
•Lowest energy (eV)  Highest eV
•Velocity is the same throughout = c = 300 000 km/s
Waves: f = Hz; λ = wavelength
Differences between Gravitational and Electromagnetic
radiation
There are two principal differences between gravity and
electromagnetism, each with its own set of consequences for the
nature and information content of its radiation, as described.
• Gravity is a weak force, but has only one sign of charge.
Electromagnetism is much stronger, but comes in two opposing
signs of charge.
This is the most significant difference between gravity and
electromagnetism, and is the main reason why we perceive these
two phenomena so differently.
• Significant Gravitational fields are generated by accumulating bulk
concentrations of matter. Electromagnetic fields are generated by
slight imbalances caused by small (often microscopic) separations
of charge.
• Gravitational charge is equivalent to inertia.
Electromagnetic charge is unrelated to inertia.
SOUND Equations
•
•
•
•
Vsound = 331.5 + .61 (Co)
v = d/t
d = vt
t = d/v
If no temp given, assume 343 m/s
 Denser the material, faster the sound!
•
•
•
•
f
λ
v
v
=v/λ
In Hertz (Hz)
= v/f
= fλ
= λ / T (period)
• Intensity (I) = Power (P) / Area (A)
– Intensity (I) = P / 4 π r2
In Watts / meter2
• Power = I (4 π r2)
In Watts
• Doppler Effect = fo = (v + vd / v + vs) fs
Light, Mirrors and Lenses
Convex Mirror
Concave Mirror
Concave Mirror
Electricity and Magnetism
Magnitude of Charge
•
•
•
•
•
•
Coulomb’s Law
FElectric = K q q’
r2
K = 8.988 x 109 Nm2/c2
q and q’ are charges of objects
r is distance between objects
Coulomb (c) and Amperage (I)
•
•
•
•
•
Amount of charge flowing through a wire in 1 second with a current of 1 ampere
Ampere is 1 Coulomb per second, the intensity (I) of the electrical current
Based on the charge of an electron
1 coulomb = 6.242 x 1018 e– Current (I) = Q / t in amperes
• Measuring the intensity of the electric current
Charge of an electron (e-) = 1.60218 x 10-19 c = 1 eV
Potential Difference (V)
•
•
•
Amount of work in an electric field to take the charge of 1 coulomb from one point to another
Volt is the potential difference across a conductor that carries a current of 1 amp
V = W/Q
– V is potential difference in Volts
• One volt = J/c
– W is work done in Joules
– Q is charge in Coulombs
Resistance (R)
• Measured in Ohms  Ω
• R = V/I
– I ohm (Ω) = 1 V / Amp
• Ohm’s Law  V = I R
– Voltage = Current in Amps x Resistance in Ohms
• Resistance in Series
– R1 + R2 + R3 + …. = RTotal
• Resistance in Parallel
– 1/R1 + 1/R2 + 1/R3 + …. = 1/RTotal
Capacitance (C)
• C = Q/V
• Measured in farads (1 coulomb per volt)
• Parallel Capacitance
– C1 + C2 + C3 + … = CTotal
• Series Capacitance
– 1/C1 + 1/C2 + 1/C3 + … = 1/CTotal
Work and Power
• Work (WE) = q V
– In Joules
• Power (P) = V q / t
– In Watts (J/s)
– Power also = V I = I2 R = V2 / R
Magnetism
• Based on charges of atom’s particles
• It is a field force – line go from N to S (Faraday Lines)
– Measured in Teslas or Gauss (1T = 100000G);
– Earth = .0001T
• All magnets have two poles – if cut it makes new poles!
• Can lose magnetism if it is heated past material’s “Curie
Temperature” and it returns when cooled
• Types of Magnetism:
– Diamagnetic: no magnetism in material
– Paramagnetic: magnetic only when in a magnetic field
– Ferromagnetic: due to e- sea model of metal, it can be
permanently magnetized