Percent Error= (Your Value-Literature Value)/Literature Value X 100 (Units Are in %)

SOL Items

Density formula D=m/v so for v= m/D
Volume= L x W x H= units3
Area formula: L x W= units2
mass: m= D x v
K=oC+273.16
oC: K-273.16= oC
Percent Error= (your value-literature value)/literature value x 100 (Units are in %)
Energy Conversion of calories to joules: 0.2390 calories (cal) = 1 Joule (J) and 1 cal=4.184 J
Atmosphere to Pascal: 1atm=101,305 Pa
Physical Properties to remember: mass, length, volume,color, density, malleability, ductility, and conductivity, crystalline shape, melting point, boiling point, refractive index.
Accuracy vs. Precision: Accuracy-Refers to how close the measurement is to the actual value while Precision refers to how close a set of measurements is together whether or not the measurements are correct
Separation of Mixtures:
Distillation-occurs when a liquid is boiled to produce a vapor that is then condensed again to a liquid. This causes the solid substances that were originally dissolved to stay in the original container and the water to go into a second receiving container.
Chromatography-Involves a solid (stationary phase) and a liquid or gas (mobile phase). The separation occurs because the liquid or gas has a faster rate than the solid.
Paper Chromatography-paper (solid) and a liquid are involved. The liquid travels up the paper and separates according to the heaviness of the individual parts of the liquid

13.Antoine Lavoisier a French Chemist (1743-1794)

Proposed the Law of Conservationof Mass: in ordinary chemical reactions, matter can be changed in many ways, but it cannot be created or destroyed.

Find on Periodic table: Atomic Number and Atomic Mass, and figure out Neutron #, Electron # and charge is negative, and Proton # and charge is positive.

Atomic Mass (Symbol Z) –Atomic number (Symbol A)=neutrons Z-A=Neutrons (neutrons have no charge and are found in the nucleus with the protons)
Note: Atomic number + neutrons =Atomic mass

Average Atomic Mass

Percent (in decimal form) times Atomic Mass for each one and then add the total

Rutherford's Gold Foil Experiment:
That the atom is mostly empty space
And that the nucleus is positive charged (because of protons) and contains almost all of the mass of the atom.
Alpha radiation is radiation that was deflected toward the negatively charged plate alpha radiation.

Made up of 2 alpha () particles

Each alpha particle contains 2 protons and 2 neutrons

Has a 2+ charge

Has a mass of 4 amu

Ex.: Ra- 88 (Radium-226) 86 Rn (radon-222) + 4 He (alpha particle) (Exact model will be shown in class)

Beta radiation is radiation deflected toward the positively charged plate beta radiation.

Consist of fast moving electrons known as beta () particles.

Each beta particle contains an electron with a -1 charge.

Ex.: C-14/6 14N/7 + 0/+1e

Half Life Formula

Amount Remaining= Original Amount of parent ÷ 2n n=half-life

Electron Configuration-The arrangement of electrons in an atom. The order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s

Rules:

Octet law: only a maximum of 8 electrons on the outermost shell (2 in the 1st level)-known as the valance electron number
The aufbau principle states that each electron occupies the lowest energy orbital available

In order of increasing energy, the sequence of energy sublevels within a principal energy level is s (2 e-), p (6 e-), d (10 e-), and f (14 e-).

The Pauli exclusion principal states that a maximum to 2 electrons may occupy a single atomic orbital, but only if the electron has opposite spins.

Represented as ↑↓

Periodic Trends:
Atomic Radius-Deals with the size of an atom
It decreases as it moves across a period
Increases as it moves down the Group
Electronegativity-The attraction it has to bond with other elements- F has the highest
It decreases as it moves down a group and increases as it moves across a period
Periodic table/Element items:

Dmitri Mendeleev put together the 1st periodic table-Table not completely correct

A period in the periodic table is all the elements in a horizontal row.

A group in the periodic table is all the elements located in the same vertical column, which is assigned a number from 1-18.

METALS

Usually shiny when smooth and clean

Conduct heat and electricity well
Solid at room temperature
Most are also ductile and malleable (meaning they can be pounded into thin sheets and drawn into wire.)
Chemically reactive
Positively charged
These atoms have only a few electrons in the outer level.
Have a tendency to lose their electrons in the outermost level.

Alkali metals: Group 1 except H

They react with water
Easily lose a valence electron and form an ion with a +1 charge.
Ends in sublevel s1

Alkaline Earth metals: Group 2

Reactivity similar to Alkali metals but not as great
+2 charge
Ends in sublevel s2

Aluminum Group (Sometimes called the Boron Group)

Group 13
+3 charge

Transition Elements for divided into 2 set:

Transition metals

Any element in columns 3-12
Has 2 electrons in the outer level - 4s2
Elements with #'s 22-28 also have a 3d sublevel
In all groups except 12, the dorbitals are only partially filled.
Share properties such as electrical conductivity, luster, and malleability with other metals.
They have magnetism

Inner transition metals: Lanthanoid Series and Actinoid Series

They are the highest energy electrons (f electrons) are inside the d sublevel and the outer level.
Both have outer shells consisting of an s2 sublevel

Lanthanoid Series- Lanthanum (57) to Ytterbium (70)

Electrons are added to the 4f sublevel instead of the sixth or outer level
Are silvery metals with relatively high melting points
Used extensively as phosphors, substances that emits light when struck by electrons.

Actinoid Series-Actinium (89) to Nobelium (102)

This series have an increasing # of electrons in the 5f sublevel
They are all radioactive.

NONMETALS

These are usually gases or brittle solids at room temperature.
Dull appearance
Insulators
Outer electrons are held closely by the nucleus
Form negative ions (anions)
Have 5 or more electrons in the outer level than metals.
They often gain electrons or share their electrons in the outermost level.

Group 14: Carbon Group

Allotropes are found in this group: forms of an element in the same physical state that have different structures and properties.

Ex. Carbon in the form of coal, Diamonds and graphite

Silicates are silicon compounds bound to Oxygen, and each Si atom is surrounded by 4 O atoms.

Group 15: The Nitrogen and Phosphorus Group

There are nonmetals (N and P), metalloids (As and Sb), and metals (Bi)
Each has 5 valence electrons and have many different properties
Charge is -3

Group 16: The Oxygen Group or Chalogens

Have 2 allotropes: Ozone, and O2
Some are oxides known as amphoteric: Those that can produce either acidic or basic solutions. Ex. Sulfur compounds like sulfuric acid (H2SO4)
Charge is a -2

The highly reactive Group 17 elements are called Halogens

Fluorine is the most reactive element- Highly electronegativity
Halogens make salts
Have 7 valence electrons and often tend to share one electron or gain one.
Have a 1- charge so they react with Group 1 the most

The extremely nonreactive Group 18 and is known as Noble gases.

All except Helium have 8 electrons in their outer level.

METALLOIDS

These are elements with physical and chemical properties of both metals and nonmetals.

Silicon and germanium are used a lot in making computer chips and solar cells.

Staircase elements between metals and nonmetals

Are often brittle solids

Oxidation Number

For metals

Positively charged
in Group 1-2: Same as Group number
in Groups 3-12: number varies
Ex. Hydrogen is in Group 1 and the Oxidation # is +1
Ex. Magnesium is in Group 2 and the Oxidation # is +2

For nonmetals

Negatively charged
Find valence electron number: 2nd number of group #
Formula: -8+valence electron number
Ex. Oxygen is in Group 16, -8 + 6= -2
Ex. Nitrogen is in Group 15, -8 + 5= -3
Ex. Chlorine is in Group 17, -8 + 7= -1

For Metalloids

Positively charged
Find the number the same way and nonmetals
Ex. Carbon is in Group #14 and the Oxidation # is +4

Ionic Bond Notes

Properties

Ionic Bonds are formed by a cation (positive charged metal) bonded to an anion (negative charged nonmetal)
Metal loses one or more electrons
Nonmetal gains one or more electrons

Note: Static electrical attraction is the basis for ionic bonds, because the positively charged ion

(cation) is attracted to the negatively charged ion (anion)

High Boiling and High Melting point
Forms a 3-D crystal lattice
Crystal lattice bonds are strong and take lots of energy to break the bonds
Usually poor conductors of electricity
Because they are solid and rigid the ions can’t move freely
Only good conductors if dissolved in an aqueous solution where they become electrolytes
Don’t consist of molecules
Strongly bonded
Can form a salt

Ionic Terms

Ion: charged particle
Anion: negatively charged ion
Cation: Positively charged ion
Salt: An ionic compound that forms when a metal atom or a positive radical replaces the H of an acid.
Ex. NaCl
Salts are excellent conductors of electricity because they are brittle solids that can easily be dissolved in an aqueous solution such as water

Ionic Bonding Problems/Diagrams

Na (+1) + Cl (-1) → NaCl
Ca (+2) + Cl (-1) +Cl (-1) → CaCl2

Naming Ionic Compounds

Many Ionic compounds contain polyatomic ions: ions made up of more than one atom.

Polyatomic ions exist as a unit, so never change the subscript
If you have to balance an ionic compound with a polyatomic ion then ( )and a subscript must be written
Ex. Ca (+2) and PO4 (-3) →Ca3(PO4)2 and named Calcium phosphate
Most polyatomic ions are oxyanions
Oxyanion is a polyatomic ion composed of an element, usually a nonmetal, bonded to one or more oxygen atoms
If a transitional metal and a polyatomic ion is involved
Ex. Cu (+2) and NO3 (-1) →Cu(NO3)2 and named Copper (II) nitrate
Note that transitional metals with varying oxidation numbers always have to state which atom was used in the chemical compound whether or not a polyatomic ion is used

Rules for Naming Ionic Compounds

Name the cation (metal) first and the anion (nonmetal) second
Monatomic cations use element name
Monatomic anions take their name from the root element name plus the suffix –ide
Ex. CsBr is Cesium bromide
Determine oxidation numbers of transitional metals compounds before naming to determine Roman number I-IV
Ex. Fe2O3 is Iron (III)oxide
Some transitional metals only have one charge
Cadmium: Cd+2
Zinc: Zn+2
If the compound has a polyatomic ion, simply name the ion
Ex. NH4Cl is Ammonium chloride
If the polyatomic ion has an oxyanions
The ion with more oxygen atoms is named using the root of the nonmetal plus the suffix –ate
Ex. NO3- is nitrate
Ex. ClO3- is chlorate
Ex. CO3-2 is carbonate
The ion with fewer oxygen atoms is named using the root of the nonmetal plus the suffix –ite
Ex. NO2- is nitrite

The oxyanions with the greatest number of oxygen atoms is named using the prefix per-, the root of the nonmetal, and the suffix-ate
ClO4- is perchlorate
IO4- is periodate
MnO4- is permanganate
The oxyanions with one less oxygen atom is named with the nonmetal and the suffix-ate
SO4-2 is sulfate
The oxyanions with two fewer oxygen atoms is named using the root of the nonmetal plus the suffix –ite
Ex. ClO2– is chlorite
The oxyanions with three fewer oxygen atoms is named using the prefix hypo-, the root of the nonmetal, and the suffix –ite
Ex. ClO– is hypochlorite
Polyatomic ions with 2 transitional metal atoms include a Di-prefix
Ex. H2PO4- is Dihydrogen phosphate
Ex. Cr2O7-2 is Dichromate
Some Hydrogen plus a polyatomic ion are named two ways
Ex. HSO4- can be named bisulfate or Hydrogen sulfate
Ex. HCO3- can be named bicarbonate or Hydrogen carbonate

Covalent Bonds

Properties

Nonmetal + Nonmetal (usually)
Most common type of bond
Covalent bonds form molecules
Form by sharing electrons
The sharing of one pair of electrons is a single bond (X-X)
Another name for single covalent bond is sigma bond symbolized by σ
Sigma bonds form from the overlap of a s orbital with another s orbital, a s orbital with a p orbital, or a p orbital with another p orbital
The sharing of two pairs- double bond (X=X)
Another name for multiple bonds is pi bond symbolized by π
Pi bonds form when parallel orbitals overlap to share electrons
A double covalent bond has one sigma and one pi bond
The sharing of three pairs-triple bond (XΞX)
A triple covalent bond has one sigma and 2 pi bonds
Bond polarity explains the attraction between the sharing
Nonpolar electrons are shared equally Ex. F-F (same electronegativity)
Polar electrons are not shared evenly Ex. H-F (different electronegativity)

Intramolecular Forces in Bonds Table

Force Basis of attraction
Ionic cations and anions
Covalent positive nuclei and shared electrons
Metallic metal cations and mobile electrons

Intermolecular Forces

Intramolecular forces do not account for all attractions between particles. There are forces of attraction called intermolecular forces.

They can hold together identical particles or two different types of particles
Also called van der Waals forces

3 types:

Dispersion forces
Sometimes called London dispersion forces
The force between oxygen molecules
Weak forces that result from temporary shifts in the density of electrons in electron clouds:

………

Attraction

Temporary attraction Temporary attraction

←|-←|-

Dipole –dipole: Attraction between oppositely charged regions of polar molecules
Stronger than dispersion forces
The more polar the molecule, the stronger the force
Hydrogen bonds: One special type of dipole-dipole dealing with hydrogen bonds

Very strong intermolecular force that is formed with a H end and a F, O, or N atom on the other dipole

Many physical properties of covalent molecular solids are due to intermolecular forces.

The melting and boiling points are relatively lower than Ionic (that is why salt doesn’t burn when you heat it but sugar will)
Many are gases are vaporized at room temperature
Hardness is also due to the intermolecular forces so covalent solids are soft in comparison to ionic solids

Naming Molecular Compounds: Rules for Binary Molecular Compounds are similar to that of naming Ionic compounds except the names include prefixes indicating the number of atoms in the molecule.

Numerial Prefixes

Mono-1
Di-2
Tri-3
Tetra-4
Penta-5
Hexa-6
Hepta-7
Octa-8
Nona-9
Deca-10

Exceptions:

H2O is water

NH3 is ammonia

Examples:

CO2 –Carbon dioxide
CO-Carbon monoxide
N2O4-dinitrogen tetroxide
SCl6 –Sulfur hexachloride

Naming Acids and Bases

Binary acids are acids with only two elements.

Prefix –hydro, stem of anion, and suffix –ic

Exception is HN3: Hydroazoic acid, where the root – azo is used for nitrogen.

Ternary acids are acids that contain 3 elements.

Usually no prefix is used and the suffix is –ic.

Exceptions:

One less O than the most common : no prefix and suffix used is –ous

Two less O than the most common: prefix hypo- and suffix –ous

One more O than the most common: prefix per- and suffix –ic

Ex. HClO3 is the most common: Chloric acid

HClO2 has one less O so: Chlorous acid

HClO has two less O so: Hypochlorous acid

HClO4 has one more O than most common so: Perchloric acid

Ternary bases

Arrhenius bases are composed of metallic, or positively charged ions and the negatively charged hydroxide ion. Therefore, these bases are named by adding the word hydroxide to the name of the positive ion.

Ex. Sodium hydroxide is NaOH.

Characteristics of Acids and Bases

Acids

Liquids are tart, sour, or sharp tasting
They conduct electricity (in solutions)-electrolytes
They produce H2 gas
Usually in liquid or gas form
pH is 0-6.9
Strong acid-have a low pH and completely ionized in an aqueous solution
The closer the substance’s pH is zero the stronger the acid
Weak Acid-have pH closer to 6.9 and are only slightly ionized in an aqueous solution
They react to metals-corrosive

Bases

Commonly found in solid form
Chemical formula except for NH3 has OH on the end
pH range is 7.1-14
Strong base-dissociates completely into metal ions and OH- ions in aqueous solution
The closer the substance’s pH is 14 the stronger the base
Some are not very soluble in water
Weak base-react with water to form the OH- ion and conjugate acid of the base
Some are insoluble in water while others are soluble
Slippery feel because bases react with oils in your skin-soaps and cleaning agents
Are electrolytes

Three primary theories of acids and bases

Theory / Acid definition / Base Definition
Arrhenius / Any substance that releases H+ ions in water solution / Any substance that releases OH- ions in water solution
Bronstead-Lowery / Any substance that donates a proton / Any substance that accepts a proton
Lewis / Any substance that can accept an electron pair / Any substance that can donate an electron pair

Examples:

I.)Arrhenius acid: HCl (g)→H+(aq) +Cl-(aq) Arrhenius base: NaOH (cr) →Na+ (aq) + OH-(aq)

II.)Bronstead-Lowery: HCl (g) + H2O → H3O+(aq) + Cl-(aq)

Acid +base→conjugate acid + conjugate base

Conjugate acid-is the particle formed when a base gains a H+ ion
Conjugate base-is the particle that remains when an acid has donated a H+ ion
Conjugate acid-base pair-consists of 2 substances related by the gain or loss of a single H+ ion

III.)Lewis: H3N: (Lewis base) + BF3 (Lewis acid) → H3N: BF3 (Product)

pH is a measurement of the H3O+ ion concentration of an acid or a base.

Problem Formulas: (Actual problem examples will be stated in class)

1.)pH=-log[H+]

2.)pOH=-log[OH-]

3.)pH + pOH=14

4.)[H3O+]=10-pH use antilog

5.)[OH-]=10-pH

6.)[OH-]=antilog (-pOH)

7.)Kw=[OH-] x [H+] which equals 1 x 10-14M

Titration

Titration-is a procedure used to bring a solution of a known concentration into a reaction with a solution of an unknown concentration in order to determine the unknown concentration or the quantity of the solute in the unknown.