Unit 4:
Describing Substances
The Model So Far
We’ve added that…
• Particles can change their phase or
temperature when energy is added.
– Examined:
•
•
•
•
Melting and Boiling Points
Heat Capacity (specific heat)
Heat of Fusion
Heat of Vaporization
Pure Substances vs. Mixtures
• Pure Substances:
– Elements
– Compounds
• Mixtures
– Homogeneous (solutions
– Heterogeneous
Physical Properties
• Characteristics of a sample that can be
observed without a change to the
sample’s identity.
– Melting/Boiling Points, Solubility, Density
• Physical Change - Change that doesn’t
change the identity
– Ex. Phase change, dissolving
Homogeneous Mixtures:
AKA: Solutions
• Solute - What is being dissolved (salt)
• Solvent - What is doing the dissolving
(water)
• When water is the solvent, it is an
aqueous solution.
Chemical Properties
• Properties observed only when there is
a change in the substance.
– Flammability, Reactivity
• Chemical change - the change of one or
more substance.
– Ex. change in color, gas production/smell,
formation of a precipitate
Electrolysis of water
• How do we know that water is H2O?
Compounds vs. Elements
• Elements contain one type of atom.
– Mixture of Iron and Sulfur retains properties of
each.
• Compounds contain two or more electrically
bonded atoms. May contain several different
types of atoms
– Iron(II) Sulfide does not look like either Iron or
Sulfur and is not magnetic like Iron.
– The atoms are compounded together and must
have a chemical reaction to separate them.
– Compounds have a set ratio of different atoms in
them: Water has 2 Hydrogen for each Oxygen.
Mass
• Antoine Lavoisier (1782) - Law of
Conservation of Mass through
experimentation with Magnesium. The new
magnesium substance weighed more than
the magnesium that he started with.
• Demonstration - balance, Mg strip
• Concluded that the mass came from the air
and developed the “law of conservation of
mass.”
Proust (1799)
• Using the combustion of water, Proust
calculated that the ratio of hydrogen to
oxygen was always the same. The mass
from the hydrogen was 11.2% and oxygen
88.8%. (Though a 2H to 1O ratio, mass of
oxygen atoms are much heavier than
hydrogen)
• Law of Definite Proportions - elements
that make up compounds always have a
certain proportion in the compound.
Law of Multiple Proportions:
Two elements can combine in
different mass ratios
EX. CO2 and CO
ONE WILL KILL YOU!
Dalton’s Theory (1803)
• All matter is made of atoms.
• Atoms are indestructible and cannot be
divided into smaller particles.
• Atoms of one element are exactly alike
but are different from other elements.
What are atoms?
• What are the properties of atoms? CHARGES
• How are different atoms different from each
other?
DIFFERENT MASS (NUMBER OF SUBATOMIC
PARTICLES)
Protons (+)
Neutrons (neutral)
Electrons (-)
• How do we know?
EXPERIMENTATION
Subatomic Particles
Table 2.1
Location In
Atom
Around Nucleus
Center
Center
Atoms have Charge….
• Tape Lab
http://physics.weber.edu/amiri/director/dcrfiles/electricity/pithBallS.dcr
http://phet.colorado.edu/en/simulations/category/chemistry
From the Tape Lab
• The T tape became positively charged
because electrons were transferred to
the B tape.
– The plastic side of the tape has a stronger
attraction for electrons than the sticky side.
• Overall number of charges does not
change (conservation of charge)
• Charges are a property of particles.
Crook’s Tube
QuickTime™ and a
decompressor
are needed to see this picture.
Crook’s Tube
• http://www.rsc.org/chemsoc/timeline//pages/1897.ht
ml
• Thomson’s model was that of plum pudding. Plums
were the electrons that could move within the
pudding (positive charges).
• When a charged item comes close to a neutral item,
the alike charges move to the opposite side of the
atom. Opposite charges move towards the item.
Smaller Than the Smallest - Subatomic Particles
QuickTime™ and a
decompressor
are needed to see this picture.
Tape and Foil
• During the tape lab you may have noticed
that the tape was more attracted to the foil
than the paper.
• The foil is a metal.
– Metals conduct electricity and heat because
electrons move easily through them.
– The tape polarized charges in the metal by
causing them to separate.
• Paper a non-metal.
– Non-metals do not conduct as well because
electrons are strongly attached to their atoms.
Electron Movement in Metals
• Electrons move through metals when an
electric current is added to it.
• In nonmetals, a current does not result
in electron movement.
• Electron attraction is weaker in metals
than in non-metals which is why they
can conduct electricity.
Compounds and Charge
• A chemical bond is an electrical force that
hold together particles that make up
compounds. Electrical forces are also
responsible for holding an atom together.
• Charges move more freely in metals than in
non-metals. This is a factor in how
compounds and molecules form.
– This will be evident by conductivity experiments.
Metals and Non-Metals
• Metallic elements are on the left side of
the periodic table.
• Non-metals are on the right side.
• The middle section includes transition
metals.
• Some elements can behave as a metal
or non-metal at extreme temperatures,
these are metalloids.
Electrolytes
• Electrolytes - solutions that conduct electricity
well.
– Compounds of nonmetals and metals
– Dissolves into ions (charged atoms)
– An outside electric force caused to the particles to
form a current.
• Non-Electrolytes - solutions that do not
conduct electricity.
– Molecules of non-metals
Conductivity Lab
• Logger Pro Demonstration
Molecular Compounds
• Does not conduct electricity in the solid and
liquid phase.
• The compound forms neutral molecules.
• The electrons are held tightly by the atom and
are unable to move to conduct electricity.
• Solids can melt at lower temperatures and do
not conduct electricity.
• Many molecular solids are non-conducting
when dissolving in water.
(exceptions are acids and bases)
Molecular Properties
• Molecules hold stronger bonds with their
electrons.
• The attractions between molecules are like
that of the top and the strop of paper, weakly
attracted so that their can be an uneven
distribution of charge.
• Molecules can be slightly attracted to each
other. Water is slightly attracted to other
water molecules.
Melting / Boiling Point
• Smaller molecules have weaker attractive
forces, lower melting and boiling points.
• Ex. Halogen gases form bonds with
themselves and are gases (boiling point
below room temp) or liquid (melting point
below room temp) at room temperature .
– (N2, O2, F2 , Cl2, Br2, I2, H2).
– Iodine is a solid at room temperature.
Chemical Formulas
• A chemical formula is a
representation of elements in a
compound or molecule and shows
the ratio of atoms in the formula.
Ex: MgCl2 or H2O
• A subscript indicates how many
atoms of a single element are in the
compound. Absence of a subscript
indicates just one of that atom.
Counting Atoms in Formulas
• List how many of each atom and how
many total atoms are in the following
formulas.
NaCl
CH4
MgCl2
Ba(OH)2
NaOH
NaHCO3
HCl
NaC2H3O2
Parentheses and Coefficients
• Atoms inside parentheses are bonded to each other
separately than the rest of the atoms in the formula.
• Ex: Ca(OH)2
– OH are bonded to each other and then bond to calcium. To be no
longer reactive. Two OH groups must react with the Ca atom.
– There are 5 atoms total
• Practice: (NH4)2SO4
• Coefficients show how many of an entire compound are
in a reaction. To figure the number of atoms involved,
first find how many are in the compound and then
multiply by the coefficient.
• Practice: 3 HNO3
• Counting Atoms WS
Naming Molecules
• Types of Molecules to Name
– Diatomic: seven (makes a seven, N2,
O2, F2, Cl2, Br2, I2 + H2)
• Uses the name of the element with -ide
at the end. Example: Nitride
– Binary Inorganic: Has two nonmetals.
NO2 - nitrogen dioxide
– Organic: Contain carbon.
Binary Covalent
• Write out the name of the first
nonmetals followed by the name of the
second nonmetal with its ending
changed to -ide. If in the same group,
name the element that is closer to the
bottom of the periodic table first. Ex.
SO2 = Sulfur dioxide
Number of Atoms
• Prefixes are used to indicate how many of
each atom is present. Ex: nitrogen monoxide
(NO), nitrogen dioxide (NO2)
• Mono
Hexa
• Di
Hepta
• Tri
Octa
• Tetra
Nona
• Penta
Deca
Number of Atoms Continued
• If only one atom of the first element is
listed. The prefix mono is not usually
included with the first element.
• Instead of monooxide, monoxide is
written.
• Unit 4: WS 1
Ionic Compound Structure
• A compound forms
when electrons are
transferred from one
atom to another.
• Ionic compounds
conduct electricity
when liquid but not
as a solid.
• The lattice crystal is
held together by
electrically attractions
of positive and
negative ions.
QuickTime™ and a
H.264 decompressor
are needed to see this picture.
Ionic Compounds
• Ions move while in liquid form, this allows for
ionic compounds to be conductive in
solutions.
• Cations (+ ions) move towards to the cathode
(- electrode) when in a current.
• Anions (- ions) move towards the anode (+
electrode) when in current.
• Other properties: brittle, high melting point,
symmetrical crystals.
Ionic Compounds
• Electrons attraction is weaker in metals
than non-metals, so electrons transfer
easily to form ions.
• Metals transfer their electrons to nonmetals.
• Metals become cations (+).
• Nonmetals become anions (-).
Ionic Solids
1. Ions are assumed to be charged spheres.
2. Ions try to surround themselves with
oppositely charged ions. This leads to a
predictable packing arrangement so that the
cation is large enough to allow anions to
surround it without touching one another.
(see previous simulation)
3. The cation to anion ratio is reflected by the
formula of the compound so that the
compound has a neutral charge.
Crystal Lattice in 3-D
• http://www.avogadro.co.uk/structure/che
mstruc/ionic/g-ionic.htm
• http://www.chm.davidson.edu/vce/crysta
ls/index.html
Ionic Compound
Nomenclature
• Ionic Compounds involve the transfer of
charge and are often between a metal
and nonmetal.
• Three Types to Name:
– Binary Ionic Compounds
– Polyatomic Ionic Compounds
– Transition Metal Compounds
Binary Ionic Compounds
• Contain only two elements.
– Named the metal first, then the nonmetal
with -ide.
• Ex: Oxide, Fluoride, Sulfide, Chloride,
Bromide, Iodide
Polyatomic Ion Compounds
• Polyatomic - a group of covalently
bonded elements that together have a
charge. Ex. NH4+ OH- or CO32• Naming:
– First positive ion first followed by the name
of the second ion. Do not change the
ending of the negative polyatomic ion.
– Ex: calcium and carbonate ion: calcium
carbonate
Transition Metal Compounds
• Transition Metals (groups 3-12) may have
a different charge for the same metal to
form different compounds. Metals that do
use a roman numeral in the name.
• No roman numeral in: Silver and Zinc
• Roman numeral example: Cu1+ =
copper(I), Cu2+ = copper(II)
• Ex: Iron(II) Nitrate has Fe2+ ions.
Ionic Formulas
• The charges from the metal and the nonmetal
should neutralize.
• The elements of groups 1,2,13 through 18
have predictable ionic charges in compounds.
1+, 2+, 3+, 4+/-, 3-, 2-, 1-, 0 (noble gases)
Ca2+, Cl1To neutralize, there must be two negative chlorine
ions to balance the 2+ charge of calcium.
CaCl2
WS 2
Representing Formulas
Precipitation Reaction
More Practice with Empirical
Formulas
• WS 4
Unit 4 Review
• Unit 4 Test