Chemistry Midyear Review SheetBy: Daphne Loyd

I. Introduction to Chemistry and Matter

We used Systéme International (SI) as our standardized measurement system Metric Prefixes:

Part of SI System 1. Length- unit is meter (m) = 39 inches. 2. Mass-. The amount of matter an object contains.Unit of mass is kilogram and also the

gram (1000 grams= 1 kilogram). Mass and weight are not the same. Weight is the force which gravity attracts matter.

3. Temperature- measures hotness of an “object”. Si unit is Kelvin (K). Also used“Celsius” in chemistry. Converting b/w Celsius and Kelvin scales.K = °C + 273

4. Time- unit is seconds (s). In chemistry, we also use minute, hour, day and year.5. Number of Particles- Unit is mole (mol). Mole = 6.02 x 10²³. In chemistry, we need to

know # of particles (atoms or molecules) in sample of matter to predict behavior.6. Electric current- unit is ampere (A). It measures the flow of electrical charge and is a

derived quantity that is based on electric current and time. EI unit for this is coulomb(C).7. Luminous Intensity- unit is candela (cd). Measures brightness of light and is not

important in chemistry.

Scientific Notation 1. M x 10ᶯ: m is the “mantissa” and n is the “exponent”2. Decimal to the left= positive exponent3. Decimal to the right= negative exponent4. To multiply two numbers in scientific notation, multiply mantissas and add the

exponents. Ex: (2.0 x 10⁵)(3.0 x 10¯²)= 6.0 x 10³5. To divide, divide mantissas and subtract the exponents.

Ex: (10.0 x 10⁵)/(2.0 x 10²)= (5.0 x 10³)6. To add (subtract) two numbers expressed in scientific notation, both numbers must

have the same exponent. The mantissas are then added (subtracted). Ex: (2.0 x 10³) + (3.0 x 10²) = (2.0 x 10³) + (0.30 x 10³)= 2.3 x 10³

Volume and Density 1. Volume- amount of 3-d space the object occupies. SI unit is cubic meter (m³). In

chemistry, (cm³) is used. 1/1000 of quart approximately = 1 cubic centimeter and 1 liter pprox..

2. 1 quart. 1 liter= 1000 cubic centimeters. 1 cm³ = 1 milliliter3. Density- Density= Mass/Volume. Measures “compactness” of a substance. Is a property

that depends only on the nature of a substance, not on the size of any particular sample of the substance. SI Unit of density is kilogram per cubic meter (kg/m³).In chem., densities of solids and liquids uses unit (g/cm³) and densities of gases are (g/dm³) or (g/L).

Reporting Measured Quantities 1. Accuracy- refers to how well a measurement agrees with an accepted value. The greater

the difference between accepted measurement and student-measured measurement, the better.

2. Precision- describes how well a measuring device can reproduce a measurement. Limit of precision depends on design and construction of device. A good general rule is that the limit of precision of a measuring device is equal to plus or minus one-half of its smallest division.

Measurement Errors 1. Observed Value- value based on laboratory measurements2. True Value- most probable value or accepted value based on references3. Absolute Error= I Observed value- True value I4. Percent Error= ( Measured Value – Accepted Value/ Accepted Value) x 100

Significant figures (sig figs) - the digits that are part of any valid measurement.Rules for Sig Figs:

I. All nonzero numbers are significant. Ex: 4.735= 4 sig figsII. Zeros located between nonzero numbers are significant. Ex: 1.0285= 5 sig figs

III. For numbers greater than or equal to 1, zeros located at the end of the measurement are significant only if a decimal point is present. Ex: 60.000= 5 sig figs, 60.= 2 sig figs

IV. For #’s less than 1, leading zeros are not significant; they merely indicate the size of the numbers that follow. Ex: 0.000200= 3 sig figs

Using Sig Figs in Calculations Multiply and Divide with Sig Figs-answer should contain as many sig figs as the less

precise measurement.o Ex: 2.3 cm x 7.45 cm= 17.135 cm²= 17 cm²

Add and Divide with Sig Figs-answer should contain as many decimal places as the measurement with the smaller number of decimal places.o Ex: 8.11 L + 2.476 L= 10.586 L= 10.59 L

Rounding a Number- we round by looking at one digit beyond the precision we need. If digit is less than 5, we don’t change the value of the preceding digit. If it is 5 or greater, we raise the receding digit by 1.o Ex: 127.36 grams= Round to 3 sig figs= 127 grams

Solving Problems 1. Use the Factor-Label Method (FLM) in Chemistry.

Ex 1: How many inches equal to 40. feet?

Ex 2: How many cubic centimeters are there in one cubic decimeter?

Ex 3: A substance called sulfuric acid flows out of a pipe at the rate of 2500 grams per second (2500 g/s). What is the flow rate in kilograms per hour (kg/h), of the sulfuric acid?

II. Atomic Structure and Model

Introduction to Atomic Model of Matter Atomic Model- A model gives an approximate picture of reality. In this case, an atomic model gives

a approximate picture to explain the properties of matter Atomic model’s history is long, and each revision gave more scientific information about matter and

atoms.Development of the Early Atoms of the Model Dalton Model

Dalton proposed the atomic theory:1. Each element has indivisible particles called “atoms”.2. In a element, all atoms are identical: atoms of different elements have different properties,

including mass.3. In a chemical reaction, atoms are not created, destroyed or changed into other types of

atoms; atoms are simply rearranged4. Compounds are formed when atoms of more than one element combined.5. Samples of a given compound always have the same relative numbers of atoms.

B/c of Dalton’s model, atoms are the basic building blocks of matter. Model explains why elements and compounds have fixed compositions by mass.

Model was missing the ability to provide clues about internal structure of atoms. J.J. Thomson’s Model, Cathode Rays and Electrons

Subatomic particles- these particles makes up the atoms J.J. Thomson’s experiment was:

High-voltage electricity was sent across a evacuated glass tube contains metal electrodes. Electricity produced stream of radiation with a negative electric charge and flowed from the

negative electrode (the cathode) to positive electrode (the anode). This radiation became known as cathode rays and the glass tube is called a cathode-ray

tube. T. examined cathode rays produced from many different sources and concluded that rays consist of streams of negatively charged particles called electrons.

It enabled the ratio of electron’s charge to its mass without knowing the value of either the charge or the mass.

Thomson’s model/ Plum Pudding Model: Atom has a (positively charged) jellylike mass with (negative) electrons scattered

Robert Millikan determined the charge on a electron by measuring a series of charges on irradiated oil droplets. Discovered that all of measured charges were multiples of a single number. Elementary charge- the number which is the charge on an electron. Elementary charge= 1.6022 x 10¯¹⁹

B/c of Thomson and Millikan, mass of electron is 9.10194 x 10¯³¹ kilogram Rutherford Model

His experiment was to : Bombard thin gold foils w/ nuclei of helium atoms and observe how these particles were

scattered throughout foils Alpha particles- Positively charge particles that has two protons and two neutrons Observed that most of the alpha particles passed through foil with little or no deflection, but

a small number were deflected at large angles.

Proposed these points: Most of volume of atom is empty space. Most of atom’s mass is concentrated in a dense, positively charged nucleus. Electrons are present in the space surrounding nucleus.

Model Developments

Experiments

Current View of Atomic Structure

Identifying Elements Name of element based on a person or a characteristic of that element Symbol- each element w/ a name has a symbol of 1 or 2 letters. Period Table of elements- lists elements by atomic number and symbols and is arranged so that

chemical and physical properties of elements show repeated pattern Group (vertical column) in Period Table of elements- organizes similarity of these properties and

numbered left to right, 1-181. Representative elements- groups 1, 2, and 13-18

Period ( horizontal rows)- properties of materials from left to right across a period change form metallic to nonmetallic

Atoms composed of a small, dense, positively charged nucleus surrounded by a large space occupied by electrons.

Nucleus = protons and neutrons

Empty space around nucleus = electrons Protons- + charge Neutrons- neutral or no charge Electrons- negative or - charge Protons mass = 1.67 x 10 ¯²⁴ g Unit for mass of particles = atomic mass unit (amu) Proton = 1 amu Neutron= 1 amu Electron= 1/1836 amu Each atom of a specific element has the same # protons as every other atom in same element Atomic Number- the number of protons in nucleus of an atom of an element Atomic mass number- the sum of the numbers of protons and neutrons in the nucleus and mass of

electrons is not included as it is so small and insignificant # of protons stays the same, but # of neutrons may vary. Isotopes- atoms of the same element that have different numbers of neutrons, and hence have

different mass numbers Ex: Carbon- either C-12 or it can be C-14 Atomic masses- mass that are fractional as shown on periodic table and is the average mass of all

the isotopes in a sample of the elementNeutrons, Isotopes and Mass Numbers Neutrons- add mass to an atom but don’t change atoms identity as an element Isotopes- contain same number of protons but have different numbers of neutrons Mass number- sum of protons and neutrons. Isotope is referred to by name and mass number

Ex: carbon-13 (c-13) or carbon-14 (c-14) Mass number- Atomic number = Number of neutronsMolecules Molecule- smallest identifiable sample of a substance

1. Monatomic molecules- noble gases or group 18 that exist as isolated atoms2. Polyatomic molecules- molecules containing two or more atoms3. Diatomic molecules- polyatomic molecules containing exactly two atoms

Molecular formula- represents every molecule in which symbol of element is preceded by a subscript that indicates the # of atoms present. We indicate presence of 1 atom by writing symbol without subscript 1. It indicates only # of atoms in each element

Ions When atoms combine, nuclei can’t change, but can either lose or gain electrons Ion-a electrically charged particle formed by loss or gain of electrons by an atom Ion has a charge b/c # of protons and electrons are no longer equal If there are more protons than electrons, ion is positively charged. If there are fewer protons than electrons, ion is negatively charged.

III. Bonding

Lewis Electron Dot Structures Shows how electrons are transferred or shared during a bond formation and has a chemical

symbol surrounded by its valence electrons Kernel- consists of nucleus and nonvalence e- Ions- charged particles with positive or negative charge Draw Compound Lewis Dot Structure:

1. Determine the total # of valence electrons2. Arrange atoms to show bonds b/w them. Central atom= least electronegativity. If after

arranged and toms doesn’t have an octet and has left over or unassigned e-, these will be double or triple bonds

Ex: C2H2: 8 +2=10

Metallic Bonds Have few valence electrons and low ionization energy Bods holding together are in solid and liquid phase which are very strong Have high melting and boiling points The kernel of central part of metallic atoms are arranged in fixed crystalline lattice Freely moving valence e- give good electrical and thermal conductivity Metallic bond= forms form force of attraction of mobile valence e- for atom’s + charged

kernel Are malleable and have a “sea of mobile e-“

Octet Rule Octet- the configuration of 8 valence e- and is max of e- except for h(2 e-) Octet rule- states the atoms react by gaining, losing or sharing e- to get a complete octet

Covalent Bonds Chemical bonds occur when attractive forces b/w atoms are greater than repulsive forces Covalent bond- formed when two atoms share e- to have a stable arrangement of e- Formed b/w nonmetal and nonmetal Double covalent bonds- a multiple covalent bond caused by sharing of two pairs of valence

e- Triple covalent bonds- sharing of 3 pairs of e- Electronegativity- measure of atoms’ tendency to attract bonded e- Nonpolar covalent bonds- electronegativity diff. close to 0 or electronegativity values equal

or close to 0 b/c the attraction of 2 chlorine atoms is equal and so electrons shared equally

Polar covalent bods- the unequal sharing of e- in a covalent bond where the electronegativity values are diff and difference is less than 1.7

Molecular Substance Molecule- smallest distinct particle of an element or compound formed by covalently

bonded atoms Molecular substances can be solid, liquid or gas depending on strength of forces of

attraction b/w molecules Molecules are are soft, poor conductors of heat and electricity, and have low melting and

boiling points Symmetrical molecules= nonpolar Ex: CO2= equal sharing of electrons Assymetrical molecules= polar Ex: H2O= unequal electron distribution VSEPR model= helps predict molecular geometry. States that the groups central atom in a

molecule generally want to be positioned as far apart as possible Using VSEPR model, the hybridization is also found by looking at how many steric groups

surround the central atom of a molecule

Ionic Bonding Ionic bonding- occurs between atoms with a large electronegativity difference, between a

metal and a nonmetal. Forms whe ions bond together b/c of oppositely charged particles Ionic bond is transfer of e-‘s. When metals bonds, the lose e-, become +charged, have smaller radii and get electron

configuration of noble gas When nonmetals bond, they gain e-, become –charged, have larger radii, and get electron

Config of noble gas. The largest electronegativity difference in a compound is ionic. Any compound with polyatomic ion is ionic bond

Distinguishing Bond Types Metals have high melting point except mercury Ionic compound has high melting points, while covalently bonded molecules have low

melting points Metallic bonds have good heat and electrical conductivity. Neither ionic or molecular solids

good conductors Ionic substances become conductors when melted(fused) or dissolved in aqueous solution

Metallic and ionic are ha rd, while covalent are soft.

IV. Electrons

Location of Electrons

Occupy the space surrounding the nucleus and Have a charge equal to, but opposite of a proton Mass= 1/1836 amu Symbol is Electrons found in orbitals Orbital- region where an electron is most likely to be found Orbitals in a atom form a series of energy levels in which electrons are found Each electron = distinct amount of energy that corresponds to the energy level it occupies Electrons gain and lose energy and so move to different energy levels in a unique way.

Instead of being able to absorb any amount of energy, an electron only absorbs a fixed amount of energy to allow it to move to a higher energy level

If an electron moves form one energy level to a different energy level, it must give off or absorb the energy difference b/w those two levels.

Ground state- when the atom’s electrons occupy the lowest available orbitals Electrons of a ground state atom fill the available spaces form lowest energy level to higher

levels until all electrons are accounted for Excited State- the unstable condition when electrons are subjected to stimuli such as heat,

light or electricity and so may absorb energy and temporarily move to a higher energy level

Electron Arrangement

Electron arrangement determines its chemical properties. Valence electrons- the outer electrons in the outer energy level or shell of the atom that the

chemical properties are based on 1st energy level= atom can contain 2 electrons 2nd energy level= can have 8 electrons Quantum theory- explains the chemical behavior of atoms Quantum numbers- a set of 4 numbers that describes electrons 1st #= major energy level of electron, principle quantum number Principle quantum number is same as number of the energy level that contains the electron If PQN is 2, then in 2nd energy level from nucleus Each energy level has 1 or more sublevels 1st sublevel = s sublevel 2nd= p 3rd= d 4th= f

3rd quantum number is relates to the orbitals in the sublevels and the orientations Electrons in s sublevels are found in orbitals w/ spherical shapes without sharp edges. Only 1 orbital in s orbital b/c there is only 1 way a sphere can be arranged P orbital is dumbbell shaped P orbital has 3 sublevels: pₓ, pᵧ and p At levels 3 and higher, the d orbital has 5 orbitals and f has 7 4th quantum #= relates to spin of an electron # indicates that each orbital can contain two electron spinning in opposite directions

Electron Configurations

Electron Config- the distribution of electrons in an atom knowing that electrons will occupy the lowest sublevel possible

V. Converting Units

Measurement Base UnitsVolume-amount of space it takes up Liters (L)Mass- amount of stuff in something Grams (G)Time SecondsTemperature- the average kinetic energy ⁰C, Kelvin (K)- absolute scalePressure= F/A Pascals, atmospheresDistance Meter (m), centimeters, millimeter, kilometer

Derive Measurements1. Density- D= M/V (g/L)2. Speed/Rate- Distance/Time3. Force- Mass + Acceleration

VI. Formulas and Equations

Each element has a symbol to identify it When writing the symbols of uncombined elements, almost all are written as monatomic

(without a subscript) Subscript- a number to the right and slightly below a number a symbol that tells the

number of atoms present Subscript is not written if there is only 1 atoms present Some elements exist in nature as two identical atoms covalently bonded into a diatomic

moleculeEx: O2 is a diatomic molecule.

Oxygen, hydrogen, nitrogen, and elements of group 17 usually are diatomic The symbols for elements is qualitative info Subscripts are quantitative info Molecular formulas- important when considering compounds formed from atoms sharing

electrons, may be a multiple of the empirical formula Empirical formula- represents simplest integer ratio in which atoms combine to form a

compound Ionic substances do not form distinct units or molecules, but rather an array of ions(charged

particles) Ionic formulas- indicate the ratio of ions in a compound Formulas of ionic substances are empirical formulas Covalently bonded substances form discrete units called molecules Atoms and compounds are electrically neutral ( do not have a net charge) b/c have equal

positive and negative charge Ions are not neutral and are positive or negative Ions w/ more protons is + charged, while ion w/ more electrons are - charged + ions attract – ions in a ratio producing a neutral compound Charge of ion is shown by a subscript following symbol of ion Polyatomic ion- a group of atoms covalently bonded together with a charge Parentheses are used to enclose polyatomic ions when there is more than one of the ions in

a unit of compound The subscripts written after the parentheses tells the reader how many of the ions are

present in the compoundFor ex:(NH4)3PO4 is three NH+

4 ions each containing 1 nitrogen atom and 4 hydrogen atoms, for a total of 3 nitrogen atoms and 12 hydrogen atoms. NH+

4 is a polyatomic ion and has a charge of 1+. The second part of the formula, PO3-

4 is the formula of one phosphate ion that contains one phosphorous atom and 4 oxygen atoms and has a charge of 3-.

Compounds can form by the attraction of oppositely charged ions

Monatomic or polyatomic ions attract each other in a ratio that produces a neutral compound

Coefficient- written in front of a formula, tells how many units of the formula are present and it applies to the entire formulaEx: 2H2O= 2 molecules of water = 4 hydrogen atoms and two oxygen atoms4Mg (NO3)2= 4 magnesium atoms, eight nitrogen atoms and 34 oxygen atoms

All compounds must be electrically neutral meaning that the sum of the charges must equal zero

For many elements the oxidation state is equal to the charge on the ion Compounds are neutral by having a equal number of positive and negative charges

Ex:1. Sodium ion (Na+) combines with a chloride ion (CI-) in a 1:1 ratio, so the compound

formula is NaCI which is neutral compound.2. For a combination of Mg2+ with CI-, a 1:1 ration wouldn’t work to get a neutral

compound. So you switch the ions charge to its opposing ions subscript the resulting compound formula MgCI2

3. Combining Na+ with NO3- is simply NaNO3

4. Combine calcium (Ca2+) with Nitrate ion (NO3-) as two nitrate ions are nitrate ions

are needed, enclose the nitrate ion in parentheses and write the subscript 2 after it forming Ca(NO3)2

Compounds are named according to the types of the elements that form them Ionic compounds or contain polyatomic ions are named the same, but covalent compounds

that contain nonmetals are named by a different method Name of a binary ionic compound comes form the names of the elements in the compound.

The positively charged particle, often a metallic ion, is placed first and the negative ion will end the formulaEx: Compound with sodium ion and chlorine ion = sodium chloride

Naming compounds with polytatomic ions:- When positive ion is metal, use the unmodified metal names plus the name of the

negative polyatomic ion. Ex: KNO3 is potassium nitrate- Most polyatomic ions are – charged. Ammonium (NH+

4) is important exception. In a compound w/ ammonium ion, the negative ion is a nonmetal with –ide ending. If ammonium ion is combined with another polyatomic ion, they each retain their namesEx: NH4CI is ammonium chloride NH4NO3 is ammonium nitrate

If a binary compound has two nonmetals, it s molecular substance and has no ions The order in which the elements are arranged in the formula can be determined by

considering the electronegativity values of the elements.

The element w/ lower negativity is written first. For example: A compound with carbon and oxygen. Carbon has a electronegativity of 2.6 while oxygen has a value of 3.4, so carbon is written first and compound will written ending in –ide. CO= Carbon Monoxide and CO2= carbon dioxide

If only one atom of the first element is present, the prefix mono- is used. If element name starts with a vowel, any final a or o in the prefix is not used.

For example: NO is nitrogen monoxide, while N2O4 is dinitrogen tetroxide Some metals have more than one common oxidation state

For ex: Iron can have an oxidation number of either +2 or +3. To help resolve this, there is the use of the stock system which uses roman numerals after

the name of the metalEx: Iron (II) chloride represents that the iron has am oxidation number of +2, and the formula is FeCI2.

In Iron (III) chloride that the iron has an oxidation number of +3 and the formula is FeCi3.

Bonding

Domain

Nonbonding

Electron-pair Geometry

Molecular

Bond Angle

Example

Model

s Domains Geometry

2 0 linear linear 180 Co2

3 0 Trigonal planar Trigonal Planar

120 BF3

CO32-

2 1 Trigonal planar Bent < 120 NO2-

SO2

O3

4 0 Tetrahedral Tetrahedral

109.5 CH4

CCI4

SO42-

3 1 Tetrahedral Trigonal Pyramidal

< 109.5

NH3

NF3

PCl3

2 2 Tetrahedral Bent < 109.5

H2O

CIO2

OF2

5 0 Trigonal Bipyramidal

Trigonal Bipyramida

l

90, 120 and 180

PF5

PCl5

4 1 Trigonal Bipyramidal

Seesaw 90, 120 and 180

SF4

BrF4+

3 2 Trigonal Bipyramidal

T-shaped 90 and 180

ClF3

BrF3

SeO32-

2 3 Trigonal Bipyramidal

Linear 180 XeF2

ICl2-

I3-

6 0 Octahedral Octahedral 90 and 180

SF6

SiF63-

SiF63-

5 1 Octahedral Square Pyramidal

90 and 180

BrF5

IF5

4 2 Octahedral Square Planar

90 and 180

XeF4

ICl4-