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