Transcript Chapter one

Chemistry - Final Review Guide
REMEMBER YOUR CALCULATOR on
the day of the exam. Also bring a sharpened
pencils and eraser.
Review topics and readings for the exam:
Chemical Reactions
Types of reactions
Composition
Decomposition
Single- replacement
Double- replacement
Neutralization
Balancing reactions
Solubility Rules:
Ions of metals from group IA and ammonium
ions are soluble
Chloride ions are soluble with the exception
of silver, mercury and lead
Sulfate are soluble with the exception of
metals from group IIA and lead
Chlorate are soluble
All others are insoluble (ppt)
Activity series rules:
1. Metals from the top of the series would
displace any metal below it
2. Metals from group IA and IIA (Ca, Sr and
Ba) would react with water, Mg reacts with
steam
3. Metals above hydrogen would displace it
from hydrochloric and sulfuric acid
Chemical composition
1. Mole Calculations:
Mol = mass/Mw
Mol = (# particles) / (6.02 x 10^23)
Mol = V in L/ 22.4 at STP
Mol = M x V in L
2 Calculations for molecular mass
Example: CaCO3 Mw = 40 + 12 + 3(16)
=100g/mol
Percent composition
Example: K2SO4
Find % K
%K = (2(Aw K)/ Mw) x100
Empirical formula and Molecular formula
Stoichiometry
Balancing Chemical equations
Interpreting the coefficient in the balance
equation
Mole-mole calculations
Mole-mass calculations
Mass-mass calculations
Limiting reactants
*Limiting reactant: the amount of reactant
that will limit the amount of product
*Theoretical Yield: the amount of product
calculated from the balance equation
*Actual (expected) Yield: the amount of
product produce in the experiment
* % yield= actual yield/theoretical yield x100
IV. Solids, liquids and gases
1. KMT - five postulates
a) Gas particles are far apart and can be
compressed, because there are a lot of empty
spaces
b) Gas particles are constantly moving and
exerting pressure on the side of the container
c) Vgas = Vcontainer (particles occupy only small
portion of the volume)
d) Gas particles don’t stick to each other
when they collide
e) KE = 3/2RT
2. Distribution change / temperature graph
3. Physical properties of gases - Expansion,
Fluidity, Low density, Diffusion and effusion
4. Gas pressure (Force/ area)
5. Measuring the gas pressure- barometer
and manometer
6 Physical equation for water - heating
curve
KMT: shape, volume, distance between the
particles and forces between the molecules
(dipol-dipol, London-dispersion and
hydrogen bonding)
8. Vapor pressure
a) Difference between boiling and
evaporation
b) factors affecting ability of a liquid to
evaporate (increases of surface area,
temperature and decrease of pressure)
c) Factors affecting the rate of evaporation of
two different liquids (particle attraction and
mass of particles)
V. Gas laws:
1. Boyles’s Law - T = const
P1V1 = P2V2
2. Charles’s Law - p = const
V1/T1 = V2T2
3. Gay-Lussac’s Law - V = const
P1/T1 = P2T2
4. Combined Law: P1V1/T1 = P2V2/T2
5. Avogadro’s Law: V1/n1 = V2/n2
6. Ideal gas Law: PV = nRT (R = 0.08205
atmL/molK)
7. Dalton’s Law: P total = Pgas + Pwater vapor for
collecting gas by water displacement
8. Calculations for molecular weight: Mw =
DRT/P
9. Real gases - at high p and lower
temperature
a) Molecules occupy volume
b) There are attractive forces between the
molecules
c) Van der Waal’s equation:
(P + n2a /V2) ( V-nb) = nRT
10. Phase diagram: triple point, determining
boiling and freezing point
VI. Solutions
1. Solution - a homogeneous mixture of two
or more substances in a single physical state.
a) Solute - substance that is dissolved in a
solution (the less quantity).
b) Solvent- substance that does the dissolving
in a solution (the more quantity).
2. Types of solution - gaseous, liquids and
solids (review the table from the notes).
* Explore 15-1: Solutions of Gases, liquids
and solids - in this activity you investigated
how substances of different physical states
can be mixed to produce
a solution.
* The nature of solutions
3. Solubility. Factors affecting the rate of
solution.
* Lab: Preparing solutions (heating, stirring
and grinding)
* Factors effecting the process of dissolving
of liquids (“Like dissolve like” rule), solids
and gases (temperature and pressure)
* Henry’s Law - The bends (15-3 activity
worksheet)
4. Saturated, unsaturated and
supersaturated solutions
a) Saturated- contains as much solute as can
possibly be dissolved under existing
conditions ( T and p).
b) Supersaturated - a solution that contains
more solute particles then are needed
to
form a saturated solution.
c) Unsaturated - a solution that has less then
the maximum amount of solute that can be
dissolved.
5. Concentration units:
a) Mass % = mass solute/mass solution x 100
b) Molarity (M) = mol of solute / volume of
solution in L
c) Molality (m) = mol of solute / kg of
solvent
d) Mol fraction = mola/moltotal
* See practice problems 15-2
* 15-2 Cooking divinity worksheet
6. Colligative properties
a) Freezing point depression
Tf = kf m
b) Boiling point elevation
Tb = kb m
* See: Practice problems 15-4
* 15-4 Explore - Colder then ice water - in
this activity you have discovered what
happens to the freezing point of water
when a substance is dissolved
VII. Equilibrium
1. The concept of equilibrium
In chemistry - most chemical reactions are
reversible processes. (Activity with the beads
and Equal rate - !6.1 explore - investigate a
reversible reaction.)
a. When the rate of forward reaction is = to
the rate of the reverse reaction - chemical
equilibrium is established.
b. Under the same condition, at equilibrium
the concentration of both reactants and
products remain constant.
c. At the same temperature the equilibrium
constant is constant
d. Equilibrium may be approached from
different starting points.
e. At other temperature the value of the
equilibrium constant differs.
2. The law of equilibrium (16.2 Review and
reinforcement - worksheet)
a) Equilibrium constant Keq and the reaction
quotient Qeq.
b) Collision theory -the molecules must
collide in order to react. In successful
collision the existing bonds in a molecule
are broken and new bonds are formed.
c) Factors affecting the rate of a chemical
reaction
* Collision frequency (depends of
concentration of the molecules and
temperature).
* Collision energy
* Orientation of the molecules
3. Reaction profile- shows the energy of
reactants and products during reaction
* Transition state - is the highest point on the
reaction profile where reactant and products
have the same potential energy.
* Activation energy - the energy require for
the reaction to achieve the transition
state
a) Energy profile for exothermic reaction
b) Energy profile for endothermic reaction
4. Types of chemical equilibrium
a) Homogeneous
b) Heterogeneous
5. Le Chatelier’s principle - when a
dynamic equilibrium is upset by disturbance,
the equilibrium will shift in a direction to
minimize the effect of disturbance
(See 16-3 Review and reinforcement).
a. Effect of concentration
b. Effect of temperature
c. Effect of pressure
d. Effect of catalyst
e. Shifting the equilibrium - the Haber
process
6. Solubility equilibrium - Ksp ( see 17.1
Problems)
sp
concentration of its ions in a saturated
solution, each raised to the power that is the
coefficient of that ion in the chemical
equation.
VIII. Acids and bases
1. Properties of acid and bases
* Acid -is any substance that produces
hydrogen ions in water
- changes blue litmus paper to red
- Sour taste
- pH  7
* Base- is any substance that produces OHions in water solution
- changes red litmus paper to blue
- Bitter taste
- pH  7
- feel slippery or soapy to the touch
Both acids and bases can undergo a
neutralization reaction.
Neutralization reaction is a reaction
between acid and base where the products
are salt and water.
2. Arrhenius definition
* Acid - is a substance that ionizes in water to
produce H+
* Base - is a substance that dissociate in water
to release OH- ions
3. Bronsted - Lowry definition
* Acid - is a substance that donates hydrogen
ions to any other substance (proton donor).
* Base - is any substance that accept a
hydrogen ion (proton acceptor).
4. Acids and bases strength
* Strong acid will ionize completely.
(HCl, H2SO4, HNO3, HClO4)
* Weak acid will ionize slightly. (CH3COOH,
HNO2, HCN, H3PO4, H2CO3)
HA(aq) + H2O(l)  H3O+ (aq) + A-(aq)
Ka = H3O+ A- HA
* Strong base will dissociate completely (
MeIAOH ).
* Weak base will provide relatively few ions
in solution
B (aq) + H2O (l)  BH+ (aq) + OH- (aq)
Kb = BH+OH-B
5. Acid- base titration
a) Principles of titration’s
measurement of the amount of a
solution
of known concentration that is required to
react completely with a measured amount of
a solution of unknown concentration.
* Standard solution - a solution that contains
a precisely known concentration of a
solute
* Titration curve - is used to represent pH
data
- Strong acid with the strong base (upward
graph - end point at pH =7).
- Weak acid with a strong base (upward graph
- end point at pH = 8.9)
- Weak base with strong acid (downward
graph - end point at pH = 5.8)
* Equivalent point - is the point at which
exactly enough standard solution is added to
neutralize the unknown solution.
* End point - the point at which the indicator
changes color.
b) Indicators - are weak acids or bases dyes
whose colors are sensitive to pH or
hydronium ion concentration.
6. Ionization of water - ionization constant
Kw = 1 x 10-14
if H3O+OH-- neutral
if H3O+OH- - acidic
if H3O+OH- - basic
pH concept
pH =- log H-
pOH = - log OH-
8. Buffers - are solutions that resist changes
in pH when an acid or a base is added
* It is composed of an aqueous solution of a
weak acid and one of its salts