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Chemistry – The study of matter, its properties, composition and change.
Unit #1: The Classification of Matter
Matter – anything that has mass and occupies space.
Matter
Pure Substances
Elements
Compounds
Mixtures
Homogeneous
(solutions)
Heterogeneous
Pure Substances – matter with a uniform (i.e. constant) composition (same throughout).
a) Elements – substances composed of only one kind of atom which cannot be broken
down using heat or electricity. Ex. Na, Br, O2, S8
b) Compounds – substances composed of 2 or more kinds of atoms and can be
decomposed using heat or electricity. Ex. H2O, NaCl, C12H22O11
Mixtures – mixtures of pure substances.
a) Homogeneous (i.e. solutions) – the same throughout. They are uniform and consist of
only one phase (i.e. solid, liquid or gas) Ex. tap water, air or brass.
b) Heterogeneous - not the same consistency throughout. May consist of more than one
phase (i.e. solid, liquid or gas). Ex. Salad dressing.
Physical and Chemical Changes
Physical change – substance maintains chemical properties but changes physically (ex. water
freezing, cutting of wood)
Chemical Change – new substance is formed. There is a rearrangement of atoms and/or ions
which creates a new compound. These changes are often irreversible.
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The Modern Periodic Table
Periodic Law: when elements are arranged in order of increasing atomic mass, chemical and
physical properties form patterns that repeat at regular intervals.
Families (groups) of elements
 Columns (i.e. vertical) of elements.
 Have similar chemical and physical properties.
Periods
 Rows (i.e. horizontal) of elements.
 Properties gradually change from metallic to non-metallic when moving from left to right
across a row.
Group Names
Alkali metals
 Group 1.
 Extremely reactive.
 More reactive towards the bottom of the group (i.e. francium is the most reactive).
Alkaline-earth metals
 Group 2.
 Very reactive (not as reactive as group 1 metals).
 More reactive towards the bottom of the group (radium is the most reactive).
Halogens
 Group 17.
 Very reactive non-metals.
 More reactive towards the top of the group (i.e. F2).
Noble gases
 Group 18.
 Inert – meaning these gases do not react with any other elements (or atoms of the same
element) and therefore will not be part of any compounds.
 Ex. Kr, He, Ne
Representative elements
 Elements from groups 1 and 2 and groups 13 to 18.
 Follow the periodic law closely.
Transition elements
 Elements from groups 3 to 12.
 Do not follow the periodic law closely.
Assignment
1. Read pages 65-69.
2. Answer questions #14-20, pg.69.
3. Pass in question #20, pg. 69.
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Physical vs. Chemical Change
Physical change


Change of state (i.e. solid, liquid or gas) or structure.
Substance maintains its chemical properties.
Ex. H2O (s) ↔ H2O (1) ↔ H2O (g)
Chemical change



Rearrangement of atoms or ions to create new compounds.
New chemicals (with their own unique set of properties) form.
Changes in state (i.e. solid, liquid or gas), colour, odour and energy may occur.
Ex: C3H8 (g) + 5 O2 (g) → 4 H2O (g) + 3 CO2 (g)
Thermal (i.e. heat) energy change


Exothermic – when energy is given off during a change.
Endothermic – when energy is absorbed during a reaction.
Note: the prefix exo- means to leave and the prefix endo- means to enter.
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Atomic Theory
Atom – is the smallest particle of an element that still represents the physical and chemical
properties of the element. These particles are composed of smaller particles (called subatomic
particles). All atoms are made from the same three subatomic particles; protons, electrons and
neutrons.
Atom
Nucleus - central region of an atom. Contains the protons and neutrons and the bulk of the mass
of the atom.
Orbit(s) – region(s) where electrons travel around the nucleus.
Subatomic particles - neutrons, protons and electrons (particles within the atom).
Nucleus
 protons (p +) - positive electrical charge (1+).
- mass of 1 amu (i.e. atomic mass unit).
- located in the nucleus.
 neutrons (n) - no charge.
- mass of 1 amu.
- located in the nucleus.
Note: the atomic mass unit (amu) is based on a concept called relative atomic mass. This is an
idea where the mass of a proton or neutron (i.e.1 amu) is compared to the mass of one carbon
atom. 1 amu is equal to 1/12 the mass of a carbon-12 atom.
Orbit
 electrons (e -) - negative electrical charge (1-)
- relatively no mass, 1/1836 the mass of a proton.
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
The number of protons defines an element. For example, only fluorine can have nine
protons. If an element has ten protons then it is neon.

The following rule applies to atoms: # electrons = # protons.

# protons (9 p+) + # neutrons (10 n) = atomic mass (19.00 amu).
Isotopes


more than one form of most elements exist.
they have the same number of protons (p +) and different number of neutrons (n).
Ex: carbon has 2 isotopes:
 carbon-12 (C-12) - has 6 p+ and 6 n
 carbon-14 (C-14) - has 6 p+ and 8 n
 note: 12 and 14 are the mass numbers.
 average atomic mass – average mass of all isotopes of an element. It is a weighted
average based on the percentage of each isotope of the element.
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Electron Energy Levels
 within an atom, electrons are arranged in different energy levels.
Sodium atom, Na atomic # = 11, #p+ =11, # e- = 11.
Valence level – outermost electron energy level. Electrons in the valence level are called valence
electrons.
1st level →2e2nd level →8e3rd level →8eBohr Diagram
2e8e2e12p+
Mg
magnesium atom
An atom is not stable unless it has a full valence level. An atom will either lose or gain electrons,
whichever is easier, to achieve a full valence level and become stable.
2e- → it will lose these two electrons to become stable.
8e- → this level is full and becomes the outer (i.e. valence) level.
8e2e20p+
Ca
calcium atom
#p+ = 20
#e- = 20
After the atom loses two electrons it will still have the same amount of protons. Since there are
now 18e- and 20p+ there becomes an imbalance of charges and the atom is said to have an
electrical charge. In this case the charge will be 2+.
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Ion


an ion is a charged particle.
charged atom or cluster of atoms.
Examples:
 Ca2+ → calcium ion
 OH- → hydroxide ion
 NH4+ → ammonium ion
Naming ions of atoms:
 Metals → name does not change. Mg2+ → magnesium ion.
 Non-metals → name changes. Keep first syllable and add the -ide suffix (i.e. ending).
7e2e9p+
8e2e9p+
F
F-
10e9p+
Example:
fluorine atom (keep first syllable, fluor- and drop –ine and add –ide suffix).
↓
fluoride ion
Radioisotopes
 radioactive isotopes of an element.
 some isotopes of some elements are radioactive. This means that they emit radiation.
 When a radioisotope emits radiation it will decay and may become another isotope of the
same element.
Radioisotope → decay → radiation + stable isotope
(unstable)
The three main types of radiation are:
 Alpha - (blocked by a piece of paper)
weak
 Beta
↓
 Gamma - (blocked by a meter of concrete) strong
carbon–14 → decays to become → carbon–12
(i.e. the carbon atom lost 2 neutrons)
“Carbon Dating”
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IUPAC- International Union of Pure and Applied Chemistry.
- Rules for naming, etc. governing body for the rules.
SATP- standard ambient temperature and pressure.
25oC and 100 kPa
For accuracy and consistency, elements on the periodic table are described under these
conditions.
Metalloids
-Found on staircase line on periodic table.
-Means →metal-like
-They have some properties of metals but are not metals.
Ex. carbon is a metalloid
-it is dull and brittle (non metal property).
- can conduct heat and electricity (metal property).
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Unit #2
Classifying Compounds
Compound – pure substance composed of two or more elements bonded together. Each
compound has unique properties that are different that the original elements that make it up.
Three Classes of Compounds
1. Ionic- combinations of cations/anions
metal/non-metal
Ex. sodium chloride (NaCl)
2. Molecular – non-metal / non-metal combinations.
Ex. sulfur dioxide - SO2(g)
Water → H2O
3. Intermetallic – metal / metal combination
- also known as alloys
- ex. Brass - CuZn
inter - means between
intra - means within
Ionic Compounds
- composed of ions - cations → positively charged ion (+)
Ex. Na+, Ca2+, NH4(+)
- anions → negatively charged ions (-)
Ex. S2-, F-, (PO4 ) 3Ions
(a) Monatomic ions – composed of only one atom
Ex. K+, CL-, Sr2+
(b) Polyatomic ions – clusters of atoms with a charge
- [poly~many: literally means → many atom ion]
NO3- Nitrate ion
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Phosphate ion
Ammonium ion
Ionic Compounds – composed of ions
- each ionic compound is composed of a cation (+) and anion (-).
 Why? Because of the law of charges (Ben Franklin), the oppositely charged ions will
attract.
 The name given to the attraction between ions in an ionic compound is called an ionic
bond.
 This results in the formation of crystalline structure (lattice).
Crystals
NaCl(s) → table salt
↓
This determines macroscopic structure of the crystal.
Macroscopic – can be seen with “naked eye.”
Microscopic – can bee seen with the aid of a microscope (ex. cells).
Molecular Level – too small for microscope [ex. atoms/molecules – electron microscope].
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**All ionic compounds are solids [i.e. Crystals] at STAP.
Another name for an ionic compound is salt.
“Formula Unit” – chemical formula for an ionic compound. Lowest whole number Ratio of
cations to anions.
Table salt is NaCl(s). This is 1:1 ratio.
Road salt is CaCl2(s) 1:2 ratio.
Formula Unit → how to determine.
magnesium chloride – composed of magnesium ions and chloride ions. Look
at their charges, Mg2+ Cl- Cl2+
2- = 0 net charge
Cations and anions form zero net charge
MgCl2(s)
Name = cation first, followed by anion. Do not indicate quantity of each ion in the name.
Name= magnesium chloride
calcium ions
6+
phosphide ions
6-
=0 net charge
Formula unit – Ca3 P2(s)
Name – calcium phosphide
calcium ions
nitrate ions
The formula unit is Ca(NO3)2(s)
Binary ionic compound → an ionic compound composed of two monatomic ions. Two kinds of
ions.
Ex. Na Cl = Na+, ClCaCl2 =
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Multi-Valent metals
Some of the transition metals can form more than one ion, each with its own charge. These
metals are called multi-valent metals.
Ex. Iron can form two different ions.
Fe3+ → iron (III) ion
Fe2+ → iron (II) ion
When an element forms an ion and it is multi-valent, we must indicate the ion charge. We use six
Roman numerals.
Roman Numerals
1-I, 2-II, 3-III, 4-IV, 5-V, 6-VI, 7-VII, 8-VIII, 9-IX, 10-X
Determine “Formula Unit” for the following:
Copper (I) oxide
Cu+ O2Cu+
.
2+ 2- =0net charge
Copper(II)oxide
Cu2+ O22+ 2-= 0net charge
CuO(S)
Copper (II) oxide
Naming
NiO(s)
Ni2+ O2nickel (II) oxide
Cr2O3(s) → chromium (III) oxide
(cation) (anion)
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Ionic Hydrates
Ionic compounds that decompose at low temperatures to produce water and an ionic
compound.
Ionic compounds with water (i.e. H2O) molecules associated with them.
Hydrous – means without water attached. Ex. – CuSO4•5H2O(s)
Anhydrous – means without water.
Ex. – CuSO4(s)
a, an}without
Ex. CuSO4•tH2O(s)
Copper (II) sulfate pentahydrate
Or
Copper (II) sulfate-5-water
Prefixes
1 – mono
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 – hepta
8 – octa
9 – ennea
10 – deca
Name Hydrates
1. Name ionic compound and quantity of water molecules.
Ex. MgSO4•7H29
magnesium sulfate heptahydrate
or
magnesium sulfate – 7 – water
Mg2+
SO42Applications
1. Desiccants – dried out
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Molecular Compounds

Non-metal / non-metal combination → ex. CO2 – carbon dioxide
ex. NO – nitrogen monoxide

Can be solid, liquid, or gas @ STAP

Do not conduct electricity when dissolved in water.
- non-metal atoms share electrons to achieve a stable state (instead of gaining or losing
electrons.)
- This “sharing” holds the atoms together is a group called a molecule.
- A molecular formula indicates the number of each in a molecule.
Ex. H2O → 2-H
CO2 → 1-C
1-O
2-O
- There are both – molecular elements (elements that form groups of like kind).
- molecular compounds (more than one element).
- Table 3.6 (pg.116) → must know by tomorrow!
- If not on table use the prefix system. → Used only for Binary Molecular compounds.
(i.e. Molecular compounds with only 2 different atoms.)
*Don’t use the prefix system on hydrogen (exception).
Steps 1. Use prefix system to indicate number of first kind of atom.
2. Use prefix system to indicate the number of 2nd kind of atom.
Ex. P4O10 → tetraphosphorus pentaoxide
Iodine heptafluoride → IF7
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Molecular Elements
Diatomic – occur in groups of two.
Elements ending in gen and halogens – H2, N2, O2, F2,Cl2, etc.

Special – P4, S8, O3 → ozone

Monatomic → all the others.
Ex. Na, Cr, Ag, etc.
See P.114 table 3.5 for extra clarification.
Acids and Bases
Acids and Bases – are corrosive solutions that occur when certain ionic and molecular
compounds are dissolved in water. (i.e. aqueous solutions).
pH scale (measure of acid or base)
Uses for pH: soil → farming, gardening. Water→ pool, aquaculture, aquarium
Empirical → fact (can test)
Theoretical → theory
Acids → empirically → turn blue litmus red
↓
Species → lichen [symbiotic relationship between
a fungus and algae.]
Theory → an aqueous solution of some molecular compounds containing hydrogen. [look
like ionic compounds] start with H________ or end with __________COOH
Ex. HCl(aq) name like an ionic compound. Start name with aqueous.
(IUPAC name) aqueous hydrogen chloride [hydrochloric acid, muriatic acid]
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H2SO4(aq) → aqueous hydrogen sulphate [sulfuric acid → car battery acid]
CH3COOH(aq) → aqueous hydrogen acetate [acetic acid, vinegar – 5to7%]
(name the positive part first)
C6H5COOH(aq) → aqueous hydrogen benzoate
Bases → empirically → turn red litmus blue
Theory → aqueous ionic compound containing a hydroxide (OH-) ion.
Ex. NaOH(aq) → aqueous sodium hydroxide (lye)
aqueous magnesium hydroxide Mg2+ OHMg(OH)2(aq)
Bases → stomach antacids: Rolaids, Tums, Diovol and Gaviscon.
Compound Classification Flow Chart
[H____/____COOH]→(yes)→acid
↓
No
↓
↓←[ionic/molecular]→↓
yes
[neutral molecular]
↓
↓←no←[does it end in OH]→ yes→↓
[neutral ionic]
[base]
Theory
Ionic
Molecular
Acid
cation/anion
m/nm
nm/nm/nm
nm/nm
Base
cation/OH
Empirical
SATP state
S
Conductivity of
aqueous solution
Litmus test
Yes
No change
s,l,g
aq
NO
No change
Yes(because act B→R
like ionics
Yes
R→B
aq
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Unit #3 Reactions
Chapter 4
3 Types of Change in Matter:
Physical, chemical, and nuclear
Physical Change- any changes where the fundamental particles remain unchanged at
molecular level. There is no change in the written formula.
Ex. evaporation, condensation, melting, freezing, cutting, etc.
Empirical
-
S ←→ L←→G
No new substance
Small energy change
Theoretical
-H2O(s) ←→ H2O(l) ←→H2O(g)
-No new molecules
- Intermolecular forces are broken and made
Chemical change –involves a kind of change in the chemical bonds within the fundamental
particles (atoms, and/or ions) of a substance. There is a change in the written formula.
Empirical
-
colour, odour, state and/or energy change
new substance formed
new permanent properties
medium energy change
Theoretical
-
2 H2O(l) + energy → 2 H2(g) + O2(g)
Atoms/ions rearranged
Chemical bonds broken and made
Nuclear change – creates entirely new atomic particles (takes place in the nucleus). There is a
new formula showing new atomic symbols.
Empirical
-
Theoretical
-
often radiation emitted
new elements formed
enormous energy change
H+H → He
New atoms formed
Nuclear bonds broken and made.
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Kinetic Molecular Theory
The central idea of the K.M.T. is that the smallest particles of a substance (atoms, ions or
molecules) are in continuous motion. As they move about, they collide with each other and other
objects in their path. The faster they motion, the greater the kinetic energy.
Energy of motion – kinetic energy
Particle Motion
Translational Motion – the motion of particles in a straight line
Rotational Motion – the spinning and turning motion of a particle
Vibration Motion – the oscillating or back-and-fourth motion of a particle
State
Solids
Empirical Properties
- definite shape & volume
- virtually incompressible
- do not flow readily
Molecular Properties
- mainly vibrational
Liquids
- assume shape of containers/
have definite volume
- virtually incompressible
- flow readily
- some vibrational,
rotational
& or transactional
Gases
- assume shape & volume
of container
- highly compressible
- flow readily
- mainly translational
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Water has the same compressibility as steel.
Gas molecules move at about the same speed as a bullet.
Volume – 3 dimensional space.
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Chemical Reactions
Evidence
Colour change
Odour Change
State Change
Energy Change
Description and Example
The final product(s) may have a different colour than the starting materials. Ex.
Bleach + pants = different colour
The final material(s) may have a different odour than the odours of the starting
material(s). Ex. sodium acetate+ hydrochloric acid→ mixture of smells like
vinegar.
The final materials may include a substance in a state that differs from the
starting material(s). Ex. silver nitrate + H2O → cloudy mixture
When a chemical reaction occurs, energy in the form of heat, light, sound or
electricity is absorbed (endothermic) or released (exothermic), mostly in the
form of heat. Ex. combustion (burning) of fuel.
Collision-Reaction Theory
- particles of the reactants must collide before any rearrangement of atoms or ions occurs.
- a certain minimum energy is required of the colliding particles, and...
- a certain orientation is required of the colliding particles for a successful rearrangement of
atoms or ions.
Law of Conservation of Mass
- In chemical changes, the total mass of matter present before the change is always the same as
the total mass present after the change, no matter how different they appear (**matter cannot be
created or destroyed). There is a rearrangement of atoms.
Mass of reactants = mass of products
(this applies to a closed system → a system in which no particles enter or leave.)
Chemical reaction
Reactants →chemical reaction →products
Balanced Chemical Equation
Is a tool that sets up an equation so that it follows the Law of Conservation of Mass. This gets
us number of each atom (reactants) = #each atom/ion (products)
Mass reactants = mass products
Balancing Chemical Reaction Equations
1. Write chemical formula for reactants and products.
2. Try balancing atom (or ion if convenient) that is most abundant.
3. Hop back and fourth across reaction arrow and balance remaining atoms/ions.
4. Check atom/ion tally for both sides.
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Classifying Chemical Reactions
1.
2.
3.
4.
5.
Formation
Simple Decomposition
Combustion
Single Replacement
Double Replacement
1. Formation → element + element + element → compound
Combining elements to form a compound.
E+E→C
Ex. Na(s) + Cl2(g) →
Na+
Cl2Na(s) + Cl2(g) → 2NaCl(s) (This is a formation reaction)
Ex. H2(g) + O2(g) →H2O(g)
2H2(g) + O2(g) →2H2O(g) (This is a formation reaction)
2. Decomposition → compound breaks down into its elements.
Compound → element + element
C→E+E
2H2O(l) → 2H2(g) + O2(g) (This is decomposition)
2NaHCO3 → 2Na(s) + H2(g) + 2C(s) + 3O2(g) (This is simple decomposition)
3. Combustion Reactions → a combustion reaction is the burning of a substance with oxygen to
produce the most common oxides of the elements making up the substance being burned.
- Exothermic (give off energy: heat, light, sound)
Technological connections → engines: transportation, autos, trains, aircraft. Heating, furnaces:
oil, kerosene natural gas. Power generation: coal, oil, gas.
Fossil fuels – oils, kerosene, gasoline, diesel, aviation fuel, coal.
Fossilized organic matter – alive at one time, compressed deep within the crust of the Earth.
Complete combustion reactions give off as products the most common oxides:
If the reactant have the following elements then common products are found to be;
C→ CO2
H→H2O
S→SO2
N→NO2
A metal →most common oxide( using the most common ion for the metal) [Ex. Pb – PbO(s)]
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Ex. a lighter (butane-C4H10)
C4H10(g) + O2 → CO2(g) + H2O(g)
C4H10(g) + 6.5O2 → 4CO2(g) + 5H2O(g)
2[C4H10(g) + O2 → CO2(g) + H2O(g)]
2C4H10 + 13 O2 → 8 CO2 + 10 H2O
Ex. burning sugar. Write a complete equation for this reaction.
C12H22O11 + O(g) →CO2(g) + H2O(g)
C12H22O11 + O(g) →12CO2(g) + 11H2O(g)
4. Single Replacement Reactions – often occur in a solution.
E+C→E+C
Ex. Na(s) + KCl(aq) → K(s) + NaCl(aq)
5. Double Replacement Reactions C+C→C+C
Ex. NaCl(aq) + KBr(aq) →NaBr(aq) + KCl(aq)
Ex. Washing Soda + Road Salt →
Na2CO3(aq) + CaCl2(aq) → NaCl(aq) + CaCO3(S)
↓
This is a precipitate.
Precipitate – is a low solubility compound formed during a solution reaction.
Double replacement special case
HCl(aq) + NaOH(aq) → HOH(l) NaCl(aq)
Acid + Base → Water + Salt
Neutralization
2HCl(aq) + Mg(OH)2(s) → 2HOH(l) + MgCl2(aq)
Acid + Base → Water + Salt
Neutralization
H2SO4(aq) + Ca(OH)2(s) →2HOH(l) + CaSO4(s)
Acid + Base → Water + Precipitate
Classifying Reactions
1. Formation E + E → C
2. Simple decomposition C → E + E
3. Combustion compound + O2
4. Single replacement E + C → E + C
5. Double replacement C + C → C + C
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Solubility
Solubility → how well a compound will dissolve in water.
Solute = the substance dissolved
Solvent = the substance doing the dissolving, usually a liquid
Solutions are categorized as having either high or low solubility
High → dissolves well.
Low→ only a small amount will dissolve.
Elements → all have low solubility in water EXCEPT chlorine. Cl2(aq)
C(S)
S8(s)
N2(g)
Molecular Compounds table 4.6 P. 153
Solubility
Examples
High
NH3(aq), H2S(aq), H2O2(aq), CH3OH, C2H5OH(aq), C12H22O11(aq),
C6H12O6(aq)
Low
CH4(g), C3H8(g), C8H18(l)
Ionic compounds→ table “solubility of ionic compound @SATP-generalizations”
a) NaCl(aq) – high solubility
b) Ba SO4 – low solubility
c) CaCO3(s) – low solubility
d) NH4CN(aq) – high solubility
e) H3BO3(aq) – high solubility
In a chemical reaction, if a new substance formed has low solubility it will become a solid and
form a precipitate.
Precipitate – is a low solubility compound formed during a solution reaction.
Example: How can you predict the SATP state of products?
Fe(NO3)3(aq) + Na3PO4 (aq) → NaNO3 (aq) + FePO4(s)
↑
↑
High solubility low solubility ( forms a precipitate)
Use the solubility table and Table 4.6 to predict the solubility of these compounds.
1.
2.
3.
4.
5.
6.
7.
8.
9.
KCl
CaSO4
LiOH
Na2CO3
FeCO3
CuCl
CH3OH
O2
K2SO4
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Unit #4
Quantitative Analysis of Reactions
The math behind chemical reactions.
-determine masses, volumes of reactants and products.
Topics: moles, molar mass, mass-mole conversions, gravimetric stoichiometry
Words that mean numbers:
Pair – 2
Dozen – 12
Baker’s dozen – 13
Ream – 500
Mole – 6.03x1023
Mole – Avogadro’s number. When dealing with small particles we need a lot of them before we
have a visible or measurable amount. We use the “mole” for this purpose.
6.02x1023 = 602, 000, 000, 000, 000, 000, 000, 000
1 mole of water molecules has a mass of 18.02 g and a volume of 18.02 ml.
Molar mass - the mass (g) of one mole of a substance.
Ex. MM H2O = 18.02 g/mol
Molar mass of octane:
C8H18
C-12.01x8= 96.08
H- 1.01x18=18.18
114.26g/mol
MM C8H18= 114.26g/mol
MM O2 = 32.00g/mol
Molar mass Conversions
Purpose – if you have a quantity of a substance, mass, you can determine the number is moles
of the material.
n= m
MM
n→ number of moles (mol)
m→ mass(g)
MM→ molar mass (g/mol)
What it the mass in grams of 2 mol of sodium chloride?
25
n=2
MM=58.44g/mol
m=(MM)(n)
m=(58.44/mol)(2mol) m=116.88g
Gravimetric Stoichiometry
Purpose: to determine masses of products and reactants in a chemical reaction.
Steps: 1. Write a balanced chemical equation.
2. Convert mass to moles
3. Mole to mole ratio (note: can only compare moles)
4. Convert moles to mass
Ex. You burn 20g of methane. How many grams of water are produced?
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
1:2:1:2
What Have CH4
m=20g
MMCH4=16.05g/mol
Mole Ratio
NH2O = 2
NCH4 = 1
NH2O = 2
1.246
Want to Get H2O
m=?
NH2O= 2.492mol
1
N=mm
MM
NH2O = (2)(1.246)
N=20g
16.05g/mol
NCH4=1.246mol
NH2O=2.492mol
MMH2O = 18.02g/mol
(1)
M=(n)(mm)
M=(2.492)(18.02)
M=44.91g of H2O
26
Unit 5: Solutions
Solutions – are homogeneous mixtures of substances composed of a solvent and at least one
solute.
Solvent – substance that does the dissolving. Usually the more abundant
substance. Ex. water, varsol, gasoline, acetone, alcohol.
Solute – substance being dissolved.
Most solutions we will be dealing with are a solute dissolved in water as a solvent. These are
known as aqueous solution. (ie. aq).
Ex. salt solution
salt – solute
water solvent
this solution is a.k.a. brine [the ocean]
Ex. sugar solution
sugar – solute
water – solvent
most soda pop
Ex. air, solution of gases [N2, O2, CO2, CH4, SO2]
Electrolytes – aqueous solutions that conduct electricity.
- ions present in solution.
- car batteries
- sports drinks
1. Separation – when molecular compounds dissolve and their neutrally charged molecules
disperse in the solvent. (ie. non-electrolytes)
Ex. sucrose dissolving in water.
Sugar solution
Sugar crystal
Separation formula
C12H22O11(s) → C12H22O11(aq)
2. Dissociation – ionic compounds dissolve in H2O.
- electrolytes, do conduct electricity.
- when ionic compounds dissolve in water, the solid compound (crystal)
breaks apart into individual ions. This process is called dissociation.
27
Salt crystal
NaCl(s) → Na+*aq) + Cl-(aq)
Dissociation Equation
3. Ionization – acids in solution
- acids are molecular compounds. They, however, ionize (atoms become
ions) then will dissociate just like an ionic compound (electrolyte).
Ex. HCl(g) → H+(aq) + Cl-(aq)
Ionization equation
H2SO4 → H+(aq) + SO42-(aq)
Svante Arrhenius Assignment
Classify A Electrolyte / non-electrolyte
B Separation / dissociation / ionization
C Write equations
Ex. KCl(s) → K+(aq) + Cl-(aq) electrolyte / dissociation
CaCl2(s) → Ca2+(aq) + 2Cl-(aq) electrolyte / dissociation
C6H12O6(s) → C6H12O6(aq) non - electrolyte / separation
CH3COOH → H+(aq) + CH3COO-(aq) ionization / electro hydrate
Acid Nomenclature (ie. acid names)
IUPAC Name
Classical Name
1. aqueous hydrogen ____ide → hydro___ic acid
ex. aqueous hydrogen chloride → hydrochloric acid
2. aqueous hydrogen ____ate → ____ic acid
ex. H2CO3(aq) aqueous hydrogen carbonate → carbonic acid
3. aqueous hydrogen ____ite → ____ous acid
ex. H2SO3(aq) aqueous hydrogen sulfite → sulfrous acid
ex. aqueous hydrogen nitrite → nitrous acid
Questions #5 & 6 Page 174
28
Identifying Ions in Solution
There are two ways to determine them, by [flame and solution] colour.

Different ions will cause solutions to appear different colours.
*Refer to back of periodic table sheet “Ion Colours” – flame colours
- solution colours

When put into a high energy situation (ex. flame) different ions will emit different
colours of light. ex. [fireworks, Northern lights] → glowing ions
Page 177, #8 & 9
Concentration of a Solution
Concentration is the numerical ratio that compares the quantity of a solute to quantity of a
solution (not solvent).
Dilute→ a small amount of solute per unit volume.
Concentrated→ a large amount of solute per unit volume (think of the can of
concentrated juice before water is added).
29
The concentration of a solution can be communicated 5 different ways:
1.
2.
3.
4.
5.
weight / weight ]
volume / volume ]% system
weight / volume ]
molar concentration ]
[parts per million –ppm ]
[parts per billion – ppb ]
[parts per trillion – ppt ]
1. Weight / of weight
amount solute
amount of solution
ex. 15% (w/w) – units doesn’t matter as long as they are consistent. (g/g, kg/kg, lb/lb, g/kg)
ex. How many grams of solute would be found in a 250g sample of 20% (w/w) solution?
% → per 100
x= (250g)(20g)
(100g)
amount of solute = amount of solute
amount of solution
amount of solution
20g = x
100g 250g
x=50g of solute
2. volume / volume (v/v) → units need to be consistent. (ml/ml, L/L, gal/gal)
ex. how many ml of acetic acid is found in a 225ml sample of 5% (v/v) vinegar?
5ml
=x
100 ml 225ml
x=(275ml)(5ml) x=11.25 mls of acetic acids
(100ml)
3. weight / volume (w/v)→expressed in (g/ml) or (kg/L)
ex. what mass of salt is found in a 350ml sample of 10% (w/v) NaCl solution?
10g = x
100 ml 305ml
x=(10g)(350ml)
(100 ml)
x=35g of solute
4. Molar concentration → how many moles of solute are present per volume of solution
mol/L, mmol/L, kmols/ML
Ex. what volume of a solution contains 0.7 mol of solute and has a molar concentration of
2.0mol/L? 1mol/L→1 M (M=molar)
2.0mol = 0.7 mols
1L
X
(2.0mol)(x) = (0.7mol)(1L)
(2.0mol)
(2.0mol)
x=0.35L of solution
Dilution – “The Solution to Pollution is Dilution”
- To weaken the concentration of a solution by adding more solvent to the
solution.
C1V2 = C2V2
30
C1 – initial concentration
V1 – initial volume
C2 – final concentration
V2 – final volume
The units of concentration must be constant on both sides. The units of volume must be the same
on both sides.
Ex. a 50ml sample of 40% (v/v) ethanol is diluted to a volume of 250ml. What is the final
solution concentration?
C1 = 40% (v/v)
V1 = 50 ml
C2 = ?
V2 = 250 ml
C1V2 = C2V2
(40%(v/v))(50ml) = (C2)(250ml)
(250ml)
(250ml)
8%(v/v) = C2
Acids and bases
According to Arrherius, bases are the result of hydroxide ions (OH-) dissolved in aqueous
solution.
Ex. NaOH(s) → Na+(aq) + OH-(aq)
↑more concentrated the more basic
According to Arrhenius, acids are the result of hydrogen ions (H+(aq)) dissolved in aqueous
solution.
Ex. HCl(g) → H+(aq) + Cl-(aq)
greater concentration of H+(aq), the more acids.
Concentration of Ions in Solution
 Sometimes it is important to be able to determine the concentration of individual ions
in solution.
 Symbol for molar concentration → [ ] Ex. [Na+(aq)] = 0.12mol/L
Steps
1. Write a dissociation or inoziation equation for the compounds.
2. Balance
3. Mole ratio
Ex. A 0.12mol/L solution of Ba(OH)2(aq). What is the [concentration] of each ion?
Ba(OH)2(s) → Ba2+(aq) + 2OH-(aq)
1(0.12mol/L) 1(0.12mol/L) 2(0.12mol/L)
0.12 mol/L
[Ba2+(aq)]=0.12mol/L [OH-(aq)]=0.24mol/L
31
Communicating Hydrogen ion, [H+(aq)], concentration
In 1909, Danish chemist Soren Sorenson came up with a system for communicating hydrogen
ion concentration. The term pH meaning “power of hydrogen” was used. The pH scale is a
method to relate the hydrogen ion concentration, [H+(aq)], in an aqueous solution.
Interesting Point – in pure water, two out of every one billion molecules will ionize to form
hydrogen and hydroxide ions.
HOH(l) → H+(aq) + OH-(aq)
Pure water is considered neutral. Hydrogen ion concentration, [H+(aq)] in neutral water is.
[H+(aq)] = 1.0X10-7 mol/L
=0.0000001mol/L
Ex. what is the hydrogen ion concentration, [H+(aq)], of vinegar? Vinegar is pH3.
pH3 [H+(aq)] = 1.0 X1-3 mol/L
= 0.001mol/L
[H+(aq)] = 1.0 X 10 -5.5 → pH 5.5
Volumetric Stoichiometry
Solution Stoichiometry
1.
2.
3.
4.
Balance Equation
Find number of moles
Mole ratio
Find volume or concentration
Ex. H2SO4 + 2KOH → K2SO4 + 2HOH
What concentration of H2SO4 solution will completely react with 15.9ml of 0.15mol/L KOH? 10
ml of H2SO4
V
C
n
KOH
15.9ml
0.15mol/L
2.385mmol
C=n/v
O=15mol/L = n/0.0159L
n=0.002385mol
n(H2SO4) = 1
n(H2SO4) = 1
nKOH
2
0.00239 2
n(H2SO4) = 0.001195
H2SO4
10ml
X
1.195mmol
32
C=n/v
C=0.001195mol/0.01L
C=0.1195mol/L
KOH
15.9ml
0.15mol/L
0.15 mol=x
L
0.0159L
(0.0159L)(0.15mol)
(1L)
NKOH=0.002385mol
Mole Ratio
NH2SO4= 1
NKOH 2
NH2SO4= 1
0.03385 2
NH2SO4 = (0.002385)(1)
2
K2SO4
NH2SO4= 0.00195mol
10ml
0.001195mol
0.01L
Concentration =
0.1195mol/L
33
Solubility Rules
(in aqueous solution)
Solids
↑Temperature
↓Temperature
↑ solubility
↓ solubility
Gases
↑Temperature
↓solubility
↓Temperature
↑solubility
Ex. trout – high oxygen demand. Cooler water.
Ex. fresh water angel fish – low oxygen demand. Warmer water.
Increase the pressure, you will increase the solubility.
Decrease pressure, decrease solubility
Ex. soda pop (dissolved CO2)
Liquids
Miscible – if a liquid will dissolve in water.
Ex. alcohols – ethanol, methanol, isopropyl
Immiscible – if a liquid will not dissolve in water.
Ex. oils
Elements
will have low solubility except for the halogens, chlorine and oxygen.
Solubility – the solubility of a solution is the concentration of a saturated solution of a
solute and a solvent at a specific temperature.
Ex. The solubility of sodium sulphate in water at 0oC is 4.76g/100ml
Saturated solution – solution at maximum concentration (is. No more solute will
dissolve).
- excess solute will not dissolve → it will settle to the bottom.
Super saturated solution – heat a solution
- therefore can be dissolve more solute
- cool solution
- solute remains dissolved beyond the saturation point for the
given temperature.
- this is know as a super-saturated solution.
34
Unit 6: Gases
Gases – fill their containers
- highly compressible
- diffuse – move spontaniously throughout any available space.
- temperature can affect either the volume m pressure of a gas or both.
Pressure → pressure = force
area
Empirical System
Pressure = lbs = psi
In2
Metric System (S.I. system internationale)
Pressure kpa (kilopascal)
1 pascal = 1N ← force[weight of one apple – 1N]
1m2 ← area
1 kPa=1000 pascals = 1000N
m2
Atmospheric pressure → barometric pressure : pressure exerted by the gasses in Earth’s
atmosphere.
One standard atmosphere→ 1atm (101.325 kPA)
Average for sea level on Earth is 101 kPa.
Standard temperature and pressure
STP → temperature of 0oC
pressure of 101.325 kPa
Standard ambient temperature and pressure
SATP → temperature of 25oC
Pressure 100 kPa
Ambient – surroundings
- room temperature
Kelvin temperature scale → has the same increments as the Celsius scale.
35
Kelvin = Celsius + 273
Ideal Gas Law
PV=NRT
P→ pressure (kPa)
V→ volume (L)
N→ number of moles
T→ temperature (K) ←Kelvin
R→ universal gas constant
8.31 (kPa)(L)
(mol)(K)
36
Unit 7 Bonding
Energy Changes (p.278 – 279)
In a chemical reaction, bonds within molecules are broken so that atoms can rearrange to
from new bonds and produce new molecules.
Ex. H2 + Cl2 → 2HCl
H – H + Cl – Cl → H – Cl + H – Cl
← → ← →
→ ← → ←
Idea: A glue holds two objects together, you must supply energy to pull objects apart. In
molecules, electrical forces act as glue.
Bond Energy → is the energy required to break a chemical bond.
Note: It is also the energy required to break a chemical bond that is formed.
 (Bonded particles) + (Energy in) → separated particles
 (Separated particles) → (Bonded Particles) + (Energy out)
Ex. Hoffman apparatus


The stronger the bond between particles, the greater the energy required to separate
them.
In a reaction, many bonds are broken and formed.
Exothermic Reaction → is a reaction that results in the release of energy in the surroundings.
(ex. energy release – heat, light and sound.)
Def: therm → heat
Exo → out
Endo → in
Ex. combustion of gasoline in a car.
2 C8H18 + 25 O2
→
18 H2O + 16 CO2 + Energy
↑
↑
↑
↓
↓
↓
Gasoline Air Combustion Exhaust
Exhaust Heat energy
↑
Chamber
Sound Energy
Pump
4
Mechanical Energy
↑
6
Light Energy
Refinery
8
↑
Oil Well
↑
Fossil Fuel
↑
Organic Matter
37
Ex. Metabolism of carbohydrates (Simple + complex sugars) in the human body.
C6H12O6 (aq) + 6 O2 (aq) → 6 H2O (aq) + 6 CO2 (aq) + Energy (heat)
(blood sugar)
(blood stream) (Blood Stream)
↑
↑
(starches/fruit sugars)
lungs
↑
↑
(plants)
air (20% O2)
Endothermic Reaction
If the overall energy changes that occur during a reaction remove energy from the
surroundings (ie. energy is absorbed) usually in the form of heat or light.
Ex. photosynthesis (in chloroplasts in cells of plants).
Energy + 6CO2(g) + 6H2O → C6H12O6(aq) + 6O2(g)
↑
↑
↑
↑
↑
(light)
air
ground
glucose
exits
Sun
leaves
roots
(starch) stomata
needles
roots
stomata
Ex. electrolysis of water
Energy + 2H2O → 2H2(g) + O2(g)
↑
Electricity
Intermolecular Bonds vs Intermolecular Forces
“within a molecule”
“between molecules”
Intermolecular Bonds → these are the covalent bonds within molecules. These bonds are very
strong as heat does not generally cause molecular substances to decompose.
(compound → element + element)
covalent bond
Intermolecular Forces → are weak forces between the molecules of a substance
→←
Observation such as surface tension, changes of state and heats of vaporization provide evidence
that there are three kinds of intermolecular forces:
1. London forces (aka dispersion forces)
2. Dipole forces
3. Hydrogen bond
38
London Forces (Dispersion Forces)
→weak attractive forces that result when electrons in one molecule are attracted by the positive
nucleus of atoms in nearby molecules.
→these forces are weak and because the distance between attracted charges are very large.
***Opposite charges attract, and same charges repel.
The boiling point of a molecular substance is an indirect measure of strength of these
London Force attractions.
Molecule
Fluorine
Chlorine
Bromine
Iodine
F2
Cl2
Br2
I2
# of electrons per
molecule
18
34
70
106
Boiling point
-188
-34.6
58.8
184
The reason why the strength of the London Forces increases as the number of electrons per
molecule increases is simple! There are more opportunities for attractive forces between
molecules.
Dipole – dipole forces
Electro negativity of an atom is a measure of how tightly an atoms nucleus will hold onto its
electrons.
The electro negativities of bonding atoms will determine the type of bond that will occur
between the atoms.
Ex. Fr= 0.7 francium loses electrons and is left positive.
F= 4.0 Fluorine takes the electrons and turns negative.
Electron negativity
Electron activity
Bond type
differences
Large
Transfer
Ionic (ionic compounds)
0.7 – 4.0
Medium
Unequal
Polar Covalent (molecular
2.1 – 3.5
compounds)
Small
Equal sharing
Covalent (molecular
3.0 – 3.0
compounds)
39
Polar Covalent Bond → is a covalent bond resulting in the unequal sharing of electrons. Since
one of the bonding atoms has a large electro negativity, it will tend to “hog” the electrons,
making that end it the molecule slightly negative (d-) overall and the other end of the molecule
slightly positive (d+) overall.
This is called “Polar Molecular”
d- d+ note: d is a Greek symbol for “delta”
ex. HCl
Dipole
* The chloride end of the HCl is slightly negative (d-) where as the Hydrogen end is slightly
positive (d+).
* Polar molecules will line up according to their positive and negative attractions.
When the oppositely charged ends of the polar molecules attract one another this is called dipole
– dipole forces. This causes the molecules to line up.
↓dipole → polar molecule
↑ dipole – dipole forces
Dipole- dipole forces are much weaker than London Forces. Also, dipole – dipole forces have
little affect on the properties of substances composed of polar molecules. However, polar solutes
will dissolve in polar solvents. Non-polar solutes will dissolve in non-polar solvents. Ie: like
dissolves like.
40
H2O
H 2.1
O 3.5
Hydrogen Bonds
When a hydrogen atom is bonded to a very electronegative atom some usual properties result.
- Fluorine (F) ~ 4.0
- Oxygen (O) ~ 3.5
- Nitrogen (N) ~ 3.0
This helped explain some properties that could not be explained by London Forces and dipole
– dipole forces.
Hydrogen Bond → are relatively strong dipole – dipole forces that occur between molecules
containing: F – H
O–H
N–H
Ex.
Very Polar molecule strong dipole – dipole forces (hydrogen bond).
Unusual properties of water that can be contributed to hydrogen bonds:
- high capacity for absorbing heat
- higher than expected boiling point (p.283)
- Powerful action of water as a solvent
- lower density of ice compared to water
Why 1 large difference between electron negativities result in highly polar covalent bonds.
41
2. The hydrogen atoms is basically stripped of its electrons making it a highly concentrated
positive pole.
*the hydrogen atom of one molecule is attracted to two “lone pairs” on the backside of the
oxygen of another molecule.
Read page 184 “Hydrogen Bonds in Biochemistry”
DNA → double helix
Structural Diagrams for Molecules
- covalent bonds
}elements N2(g)
- non-metal/non-metal }compounds H2O
Lewis Dot diagram/Electron dot diagram
- group number – last digit – work around symbol
Ex, phosphorus – group 15 – 5 valence electrons
* A valance level can hold max: 8 electrons.
* Has 4 Orbitals, each can hold 2 electrons.
* A full orbital is called a lone pair (non-bonding).
*Unpaired electron → (bonding electron).
42
How to do structural diagrams of molecules. (single bonds)
Ex. water H2O
Group 1
HH-
Group 16
Bond Capacity – the maximum number of single covalent bonds that an atom can form.
43
Molecular Shapes
1. linear – straight line
2. V-shaped
3. Pyramidal
4. Tetrahedral
H-H
O=O
O=C=O
44
Chemistry 112 – Terms

Acid – a substance that forms a conducting, aqueous solution that turns blue litmus paper
red, neutralizes bases, and reacts with active metals to form hydrogen gas.

Alkali metals – the family of elements corresponding to Group 1 of the periodic table of
elements.

Alkali – earth metals – a family of elements in Group 2 of the periodic tables of
elements

Anhydrous – the form of a substance without any water of hydration.

Anion – a historical name for a negatively charged ion.

Aqueous – a solution that has water as a solvent.

Atom – the smallest part of an element that is representative of the element; a neutral
particle composed of a nucleus containing protons and neutrons, and with the number of
electrons equal to the number of protons.

Atomic mass – historically defined as the mass of an element that combines with one
gram of oxygen.

Atomic number – a characteristic number for an element; believed to represent the
number of protons in the nucleus of an atom.

Avogadro’s number – equal volumes of gases at the same temperature and pressure
contain the same number of molecules.

Base – compound forming a conducting, aqueous solution that turns red litmus paper
blue and neutralizes acids.

Binary ionic compound – a compound that contains only two kinds of monatomic ions.

Cation – a historical name for a positively charged ion.

Chemical change – a change in which new substances with different properties are
formed, as evidenced by changes in color, energy, odor, or state.

Closed system – one in which no substance can enter or leave.

Coefficient – the number of molecules or formula units of a chemical involved in a
chemical reaction.

Combustion – the rapid reaction of a chemical with oxygen to produce oxides and heat;
complete combustion produces the most common oxides.
45

Compound – a pure substance that can be separated into elements by heat or
electricity; a substance containing atoms of more than one element in a definite fixed
proportion.

Concentrated solution – a homogenous mixture with a relatively high ratio of solute to
solution; e.g., a saturated solution.

Concentration – the ratio of the quantity of solute to the quantity of solution or solute.

Covalent bond – the simultaneous attraction of two nuclei for a shared pair of electrons.

Crystallization – the process of obtaining a solid by evaporating the solvent or cooling a
concentrated solution.

Diagnostic test – a short and specific laboratory procedure with expected evidence and
analysis used as an empirical test to detect the presence of a chemical.

Diatomic – composed of two atoms.

Dilute solution – a homogenous mixture that has relatively little solute per unit volume
of solution.

Dilution – the process of decreasing the concentration of a solution, usually by adding
more solvent.

Dissociation – the separation of an ionic compound into individual ions in a solvent.

Distillation – the process of vaporizing and then condensing a liquid.

Double replacement – the reaction of two ionic compounds in which cations and anions
rearrange, producing two new compounds.

Electrolyte – a solute that forms a solution that conducts an electric current; a substance
that ionizes in water to form individual ions.

Electron – a small, negatively charged subatomic particle; has a specific energy within
an atom.

Element – a pure substance that cannot be further decomposed chemically; composed of
only one kind of atom.

Empirical – relating to past experience or experiments.

Equilibrium – a state of a closed system in which all measurable properties are constant.

Family – a group of substances with similar properties; e.g., a family of elements or an
organic family.

Formation – the reaction of two or more elements to produce a compound.
46

Formula unit – the smallest amount of a substance that has the composition given by
the chemical formula.

Gas – a substance that fills and assumes the shape of it container, diffuses rapidly, mixes
readily with other gases, and is highly compressible.

Group – a column of elements in the modern periodic table.

Halogens – family of elements corresponding to Group 17 of the periodic table of the
elements.

Heterogeneous mixture – a non-uniform mixture consisting of more than one phase.

Homogenous mixture – a uniform mixture consisting of only one phase.

Hydrate – a compound that decomposes at a relatively low temperature to produce water
and another substance; a compound containing loosely-bonded water molecules.

Immiscible – two liquids that form separate layers instead of dissolving.

Inert – chemically non reactive.

Ion – single particle or group of particles having a net positive or negative charge.

Ionic bond – the simultaneous attraction among positive and negative ions.

Ionic compound – a pure substance formed from a metal and a nonmetal; crystalline
solid at SATP; has relatively high melting point; and conductor of electricity in molten or
aqueous states.

Ionization – the process of converting an atom or molecule to an ion.

Isotope – a variety of atoms of an element; atoms of this variety have the same number
of protons as atoms of other varieties of this element, but a different number of neutrons.

IUPAC – International Union of Pure and Applied Chemistry; the organization that
establishes the conventions used by chemists.

Litmus – a plant dye commonly used as an acid-base indicator.

Malleable – the ability to be formed or stretched by hammering or rolling.

Matter – anything that has mass and occupies space.

Metal - element that is shiny, silvery, and a flexible solid at SATP; most metals are good
conductors of heat and electricity and they tend to form positive ions.

Metalloid – element near the staircase line in the periodic table; solids with very high
melting points that are either non-conductors or semiconductors of electricity.
47

Miscible – liquids that mix in all proportions and have no maximum concentration.

Mole – SI base unit for the amount of a substance; one mole is the number of entries
corresponding to Avogadro’s number (6.02 X 1023/mol).

Molecular compound – a pure substance formed from nonmetals; solid, liquid, gas or at
SATP, relatively low melting point, and non-conducting in any state.

Molecule – a particle containing a fixed number of covalently bonded nonmetal atoms.

Multi – valent – the ability of an atom to form a variety of ions.

Neutral – an atom or molecule is considered to be neutral when its net charge is zero.

Neutralization – a double replacement reaction of an acid with a base to produce water
and a salt of the acid.
Neutrons – uncharged, subatomic particles present in the nuclei of most atoms.


Noble gases – the family of elements corresponding to Group 18 of the periodic table of
elements.

Non-electrolyte – a solute in a solution that does not conduct an electrical current; a
substance that does not produce ions in solution.

Nucleus – the central region of an atom that contains most of the mass and all of the
positive charge of the atom.

Orbital – according to the theory of quantum mechanics, a region of space where there is
a high probability of finding electrons of a particular energy; the orbital may contain a
maximum of two paired electrons.

Period – a horizontal row of elements in the periodic table whose properties change from
metallic to nonmetallic from left to right.

Periodic law – chemical and physical properties of elements repeat themselves at regular
intervals when the elements are arranged in order of increasing atomic number.

Polyatomic ion – a group of atoms with a net positive or negative charge on the whole
group.

Precipitate – a low solubility solid formed from a solution.

Pressure – force per unit area.

Protons – positively charged, subatomic particles found in the nuclei of atoms.

Pure substance – homogenous matter that has a definite set of physical and chemical
properties, that cannot be separated by physical changes; elements and compounds.
48

Radioisotope – a radioactive isotope of an element.

Salt – an ionic compound whose cation is not H+ and whose anion is not OH-; also the
common name for sodium chloride.

SATP – standard ambient temperature and pressure; 25C and 100kPa.

Saturated solution – a solution that is in equilibrium with undissolved solute and
contains the maximum amount of dissolved solute at specified conditions.

Single replacement – the reaction of an element with a compound to produce a new
compound.

Solubility – concentration of a saturated solution of a solute in a solvent.

Solute – a substance that is dissolved in a solvent.

Solution – a homogenous mixture of dissolved substances containing at least one solute
and one solvent.

Solvent – medium in which a solute is dissolved; usually the liquid component of the
solution.

Valence electrons – the electrons in the outermost energy levels of an atom.