Final Exam Review Fall 2012
Dec 26, 2015
Final Exam ReviewFall 2012
Measurement and Significant Figures Rules for Significant Figures
Any number that is NOT zero is significant.
Any zeroes between 2 numbers is significant.
Any zeroes before any numbers is NOT significant.
Any zeroes after a number is significant ONLY if there is a decimal point anywhere in that number!
Examples 43.59g
5043mL
0.000674mol
250J
670.J
Measurement and Significant Figures
ADDING/SUBTRACTING
You answer has the same # of significant figures as the number that has the LEAST AFTER THE DECIMAL
Example
456.12mL + 76.789mL = 532.909mL 532.91mL
MULTIPLYING/DIVIDING
You answer has the same # of significant figures as the number that has the LEAST TOTAL # OF SIG. FIGS.
Example
8567.10J/ (100.g × 4.184J/g°C) = 20.47586042°C 20.5°C
Measurement and Significant Figures Convert metric units
Use King Henry Died by Drinking Chocolate Milk.
K H D base d cm
Figure out how many spaces you need to move the decimal!
Examples
Convert 45.6 kg to g
How many L are in 12,980mL
Measurement and Significant Figures Converting other things – use Dimensional Analysis
Example How many inches are in 7.3 feet. (12 in = 1 ft.)?
7.3 ft | 12 in = 88in
| 1 ft
How many cups are in 45.0 ounces (oz.) [8oz. = 1 cup]
45.0 oz. | 1 cup = 5.63 cups
| 8 oz.
Measurement and Significant Figures Converting Temperatures
TF = 1.8(TC) + 32° or TC = 0.56(TF - 32 °)
Example- Convert 98.6 °F to °C.
TC = 0.56(98.6 ° - 32 °) = 37.3 °C
Converting Pressures 1 atmosphere = 760mm Hg = 760 torr = 101,325 Pa
Example- Convert 0.85atm to torr
0.85 atm | 760 torr = 646 torr
| 1 atm
Matter
Definite shape & volume
Definite volume but NOT shape
No definite shape or volume
Matter – Chemical vs. Physical Changes
Physical Changes
NEVER change what the substance is made of Examples:
Boiling
Freezing
Cutting
Grinding
Chemical Changes
ALWAYS changes into a new substance Examples:
Burning
Digesting
Rusting
Oxidizing
Reacting
Called a chemical reaction!
Matter: Pure Substances vs. Mixtures
Pure Substances
Made of ELEMENTS or COMPOUNDS chemically bonded together
Can ONLY be broken down in 2 ways Chemical reaction if it’s
a compound
Nuclear reaction if it’s an element
Mixtures
A blend of 2 or more pure substances Homogeneous – can’t
see each part (tea)
Heterogeneous – can see each part (rocky road ice cream)
Can be broken down by physical changes
Matter – Elements and the Periodic Table
Matter – Atomic Structure
Atoms make up each element
Atoms are made of protons (p+) and neutrons (n0) in the nucleus and electrons (e-) in the electron cloud surrounding the nucleus. Protons have a + charge
Neutrons have no charge
Electrons have a – charge (smallest mass)
Atomic # = #p+ (#e-)
Mass # = #p+ + #n0
Can be rounded from the atomic mass (for the most common isotopes!).
Isotope = atoms of an element that have the same #p+ BUT different #n0
Matter – Atomic Structure
Practice – How many protons, neutrons, and electrons are in:
20782Pb 41
20Ca
Matter
Electrons exist in orbitals in the electron cloud Called the ground state – they have the least amount of
energy possible
When photons of energy are added to an element, it’s called the excited state!
Nuclear Chemistry – Inside the Nucleus The ONLY way to change an atom of 1 element into
another element
Nuclear fusion – 2 atoms combine to create a larger atom + lots of energy!!!! Stars use fusion to create all the natural elements!!!!!
Nuclear fission – an atom splits into smaller particles
Nuclear particles – released in fission or fusion Alpha particles – nucleus of a helium atom
Beta particles – electrons
Gamma particles - photons
Bonding
Ionic Bonds
Occur between a metal and a nonmetal
Electrons are transferred from the metal to the nonmetal
Covalent Bonds
Occur between 2 nonmetals
Electrons are shared between the atoms
Ionic Bonds – transfer of electrons
Covalent Bonds – share electrons (Lewis Structures) 1. Add all the valence electrons
for each atom
2. Divide by 2 to determine pairs of electrons
3. Find center atom and draw end atoms around it
4. Draw a line (bonding pair of electrons) between each end atom and the center atom
5. Determine the lone pairs of electrons left over.
6. Place them around the end atoms 1st, then around the center to make sure that all atoms have 4 pairs around it.
7. If you run out of lone pairs, create double or triple bonds.
Periodic Trends
Fluorine:- Has the smallest
atomic/ionic radius- Has the largest
ionization energy- Has the largest
electronegativity
Francium:- Has the largest atomic/ionic radius- Has the smallest ionization energy- Has the smallest electronegativity
Nomenclature: Names & Formulas Writing Formulas
Type I Ionic Compounds (regular metals)
Write the symbol & charge for the cation (metal)
Write the symbol & charge for the anion (nonmetal/polyatomic ion)
Criss-cross the charges.
Example
Calcium phosphate
Ca2+ PO43-
Ca3(PO4)2
Nomenclature: Names & Formulas Writing Formulas
Type II Ionic Compounds (have transition metals)
Write the symbol & charge for the cation (metal) The charge is given by the Roman numerals!
Write the symbol & charge for the anion (nonmetal/polyatomic ion)
Criss-cross the charges.
Example
Manganese (III) nitrite
Mn3+ NO21-
Mn(NO2)3
Nomenclature: Names & Formulas
Writing Formulas Type III Covalent Compounds (2 nonmetals
Write symbol of 1st element and prefix becomes a subscript
Do the same for the 2nd element.
Example
Disulfur hexachloride
S2Cl6
Prefixes1- mono2- di3- tri4- tetra5- penta6- hexa7- hepta8- octa9- nona10- deca
Nomenclature: Names & Formulas Writing Formulas
Acids
Binary Acids – Use the prefix hydro in the name!!!!
Write H1+.
Write anion and charge
Criss-cross charges
Example
Hydrosulfuric acid
H1+
S2-
H2S
Oxyacids – don’t use any prefix
Write H1+
If name ends in –ic acid, use polyatomic ion ending in –ate.
If name ends in –ous acid, use polyatomic ion ending in –ite.
Criss-cross charges.
Example
Sulfuric acid
H1+ SO42-
H2SO4
Nitrous acid
H1+ NO21-
HNO2
Nomenclature: Names & Formulas
Writing Names Type I Ionic Compounds
Write name of cation (metal).
Write name of anion (nonmetal/polyatomic ion). If it is just a nonmetal, change ending of element to –ide.
Example
Na2O sodium oxide
KMnO4 potassium permanganate
Nomenclature: Names & Formulas
Writing names Type II Ionic Compounds (Transition metals)
Have to write the charge of the metal cation as a Roman Numeral
Write name of cation (metal).
Write the original charge of the metal as a Roman numeral.
Write the name of the anion (nonmetal/polyatomic ion). Change ending of nonmetal to ide.
Example
FeF3 iron (III) fluoride
Cu2CO3 copper (I) carbonate
Nomenclature: Names & Formulas
Writing names Type III covalent compounds
Write the name of the 1st element, change subscript to prefix. (Remember, don’t use mono for the 1st element.)
Write the name of the 2nd element, change the subscript to prefix. Change ending to –ide.
Example
SO3 sulfur trioxide
P4O10 tetraphosphorous decoxide
Nomenclature: Names & Formulas
Writing names Acids
Binary Acid (only 2 elements)
Write hydro + 2nd element’s name + ic acid
Example HCl hydrochloric acid
Oxyacids (polyatomic ion with oxygen in it)
If polyatomic ion’s name ends in –ate, change to –ic acid.
Example H2CO3 carbonate carbonic acid
If polyatomioc ion’s name ends in –ite, change to –ous acid.
Example H2SO3 sulfite sulfurous acid
The Mole
Balancing Equations
When you balance, you write COEFFICIENTS (numbers that go in front of a formula) to make sure that each side of the reaction has the same number of atoms of each element.
Example
NaCl + F2 NaF + Cl2
H3PO4 + Mg(OH)2 H2O + Mg3PO4
Types of Reactions – 5 Types
Synthesis: 2 reactants 1 product
Decomposition: 1 reactant 2 or more products
Single-replacement: 1 element + 1 compound 1 element + 1 compound
Double-replacement: 2 compounds switch ions
Combustion: 1 hydrocarbon + oxygen CO2 + H2O
Stoichiometry
Given an amount of 1 substance, how much of another substance is needed or can you make?
Uses your understanding of moles and balanced equations to solve problems.
Mole Ratio – used to convert between 1 substance and another substance.
You may have to convert your original substance to moles 1st.
You many have to convert your new substance back into mass from moles after you use the mole ratio.
Stoichiometry
KMT – 5 assumptions
1. Gases are made of tiny particles far apart relative to their size
2. Gas particles are in continuous, rapid, random motion
3. There are no attractive forces between molecules
under normal conditions of temperature and pressure
4. Collisions between gas particles and between particles and container walls are elastic collisions.
5. All gases at the same temperature have the same average kinetic energy. The energy is proportional to the absolute temperature.
KMT & n, T, V, & P
Why is the air heated in a hot air balloon to inflate it?
Attractions that come into play when gases become liquids
van der Waals forces – weak temporary attracts between 2 molecules
Hydrogen bonds – a van der Waals force that deals with the hydrogen on 1 molecule and a nonmetal on another molecule
Heat!
q = mCΔT
Endothermic (+q) – heat is added, gets hotter
Exothermic (-q) – heat is removed, gets colder
ΔE = q + w
ΔE = change in energy
q = heat
w = work
q is + (endothermic)
q is – (exothermic)
w is + when work is done on the system
w is negative when work is done by the system
Molarity (M)= a measure of concentration
M = n
V n = moles of solute
V = volume of solution in L
Acid-Base Theories
Arrhenius
Acid- has H1+ ions in it
Base- has OH1- ions in it
Brönsted-Lowry
Acid- donates H1+
Base- accepts H1+
Conjugate base- what becomes of the acid after donating H1+
Conjugate acid- what the base becomes after accepting an H1+
Neutralization of Acids & Bases
MAVA = MBVB