Honors Chemistry 2020-21

Properties of Matter

Matter: anything that has mass and occupies space

States of Matter: Solids, Liquid & Gas

1. Solids: have definite volume and shape

~Incompressible and strong chemical bonds, high density

~Crystalline structures are rigid and have an order structure

~Amorphous structures lack order

2. Liquids: have definite volume

~Not as compressible as gas

~More dense than gases

~Particles are fluid-like and tightly packed; slide across one another

3. Condensed Phases: Particles are very close together; liquids and solids

4. Gases: No definite shape or volume

~Small particles travel in constant, random motions

~Elastic collisions

~Low density

~Low attractive & repulsive forces

~Compression & Expansion

Classification of Matter

Mass vs Weight:

~Mass measures amount of matter; constant

~Weight is dependent on gravity; varies

Physical Properties & Changes

Physical Properties: Can be observed or measured without changing the composition of a substance ~Ex. Color, smell, solubility, conductivity, melting point, boiling point, density, malleability, hardness

Viscosity- the tendency to resist motion

● Greater viscosity moves slower

Conductivity- the ability to allow heat to flow

3 Types of Changes in Matter

~Physical changes in the state of matter

~Chemical changes in the substance

~Nuclear changes of the nucleus

Separation Techniques

~Distillation: separated by boiling points

~Filtration: separates liquids and solids by mass or density

~Chromatography: separates liquids in a mixture by attraction

Phase Changes

~Ex: melting, freezing, vaporization, condensation, sublimation, deposition

~Endothermic vs Exothermic

● Endothermic uses more energy than exothermic

● Endothermic is the absorption of heat; exothermic is the release of heat

Chemical Properties and Reactions

Chemical Properties: can be observed or measured while changing the composition of a substance; going from one substance to a different substance

~Ex: flammability, reactivity

Reactivity: describes how readily a substance combines chemically

Chemical changes: produce matter with a different composition

● Atoms are rearranged; chemical bonds are formed and broken apart

Precipitate: a new formed solid in a reaction

Evidence of a Reaction: heat energy is transferred, a precipitate or gas is formed, odor or color change, light energy is produced

~Ex: rusting, exploding, burning, decomposing, oxidation

Flow of Energy

Endothermic Reactions: Vaporization and sublimation

● Moves up the arrow (see visual above)

Exothermic Reactions: Condensation and deposition

● Moves down the arrow (see visual above)

Law of Conservation of Energy: Energy can not be created or destroyed

● Potential Energy: stored energy in chemical bonds

● Kinetic Energy: energy in motion

● Thermal Energy: heat energy

Heating Curve: graphs phase changes

● Plateau: flattening the curve

Scientific Notation

Coefficient Exponent

● Positive Exponent: Moves to the right

● Negative exponent: Moves to the Left

Significant Figures

● All nonzero numbers are significant

~Ex: 3.55

● Zeros between numbers are significant

~Ex: 3.05

● Zeros at the end of the number after the decimal are significant (Trailing)

~Ex: 3555.00

● Zeros before nonzero numbers are not significant

~Ex: 0.00355

● Zeros at the end without a decimal point are not significant

~Ex: 355000

● All digits shown in proper Scientific notation are always significant

Adding and Subtracting

Rule: Answers should be rounded to contain the same number as decimal places as the number with the least decimal places

● 6.78 + 3.5619 = 10.34

Multiplying and Dividing

Rule: Answer should be rounded to contain the least amount of sig figs

● 8.93 x 2.0 = 18

Dimensional Analysis

Accuracy vs Precision

Accuracy: measures how close you are to the true value

Precision: measures how close many measurements are to each other

Units & Metric DA

Gases

Constants: 1.00atms = 760.0 torr = 76.00cmHg = 1.01325 pa= 101.325kPa = 1.01325 bar = 14.696 psi

Pressure Changes

As altitude increases, pressure decreases

Temperature

Celsius, Fahrenheit, Kelvin

● K=^oC + 273.15

Boyle’s Law

(P1)(V1) = (P2)(V2)

Charles’ Law

Gay Lussac’s Law

Combined Gas Law

Kinetic Molecular Theory

~Volume of Individual Particles = 0

~Always in Motion

~Collisions of particles with container walls cause pressure exerted by gas

*elastic collisions

STP: Standard Temperature and Pressure

● Standard Temperature: C or K

● Standard Pressure: ATM

Law of Definite Proportions

Regardless of amount, a compound is always composed of the same elements in the same proportion by mass

~Compounds of different mass rations have the same elements found, but they will have different properties ● CO; CO₂

**Dalton

● Democritus: Greek Philosophers; cut up rocks

● Dalton: experimented with gases, confirmed atoms

● Mendelevium: Periodic table my mass

● Mosely: Periodic Table by atomic number

Naming Ions

~Cations: + charge

~Anions: - charge

*Transition metals are polyvalent and have different charges

● Write as Roman Numerals

----------> Cu⁺: copper (I) iron

**cation

~ cation: element name & ion

~ cation (transition metal): element (roman numeral) ion

~anion: element name & -ide

~metals: left of the staircase

~nonmetals: left of the staircase

Covalent Compounds

~nonmetal + nonmetals

~share electrons, molecules, neutral charge, low melting and boiling points, weak chemical bonds

~Name: Greek Prefix(first full name) + Greek Prefix (name + -ide)

● N₂H₄: dinitrogen tetrahydride

Ionic Compounds

~metal + nonmetals

● Metals first nonmetals second

● -ide, -ite, -ate

~transfers electrons, high melting and boiling points, strong chemical bonds, conducts electricity when liquid

~Binary Ionic: Two elements

● Aluminum Sulfide

○ Al³⁺S^2-

○ Al₂S₃

Naming Acids

Acid: compound with one or more hydrogen atoms and produces H⁺ions when dissolved ● HnX

~H: hydrogen cation

~n: subscript

~X: monatomic or polyatomic ion

The Mole

6.022 x 10²³ elementary entities in 1 mol

Molar Mass: the mass of a compound divided by whole

● part/whole

Molarity

Chemists need to make solutions with precise concentration

~The concentration of a solution tells you how much solute that is dissolved in an amount of solvent Solvent: what is dissolved

Solute: Capable of dissolving

Molecular Formula: tells the actual number of atoms in a compound

Empirical Formula: reduced version of the molecular formula

Finding Molecular Formula

● Calculate the MM of the empirical formula

● Calculate the molar mass of each atom/element

● Multiply the Subscripts

● Write the Formula

Chemical Reactions

Evidence: Heat energy, precipitate or gas, change in smell and color, formation of light energy Aqueous (aq): mixed with water

● Balance metals

● Balance nonmetals

● Balance hydrogen

● Balance oxygen

Types of Chemical Reactions

● Combination: A + B -----> AB

● Decomposition: AB ------> A + B

● Single Replacement: A + Bx -----> Ax + B

● Double replacement: Ax + BY -----> AY + BX

● Combustion: Hydrocarbons + Oxygen Gas -----> Heat + Water +Carbon Dioxide ● Acid Base: Acid + Base ----> Salt + water

Predicting Products

● Define type of reaction

● Drop and switch the charges of reactants

● Balance

*Diatomic Elements

*Hydrogen,Nitrogen,Fluorine,Oxygen,Iodine,Chlorine,Iodine,and Bromine

Molar Ratios

2Fe₂O₃ + 3C -----> 4Fe + 3CO₂

~2 mol 2Fe₂O₃:4mol Fe

*AND MORE

Types of Yield

Theoretical Yield: max amount of a product in grams that is formed if all the limiting reactant is used up in a reaction

Actual Yield: the measured amount of product actually obtained from an experimental reaction

Percent Yield: ratio between the two yields

*can be above 100%

Percent Yield= (actual yield/theoretical yield)x100

N2+ 3H2 -------> 2NH3

Error

Excess Reactant

Convert from mass of limiting reactant found to Mass of Excess Reactant Used

Atoms

Subatomic particles: protons, neutrons,and electrons

Protons: positive charge, located inside the nucleus, 1 amu

Neutrons: no charge, located inside the nucleus, 1 amu

Electrons: negative charge, flies around the nucleus in the energy levels/orbitals/ electron clouds

Energy Levels: fixed distance from the nucleus of an atom where electrons may be be found Nucleus: Center of the atom

Isotopes: atoms that have the same atomic number but different masses due to a change in the number of neutrons

Atomic Number: number of protons in an atom

**the number of electrons are equal to the number of protons

Mass Number= Protons + Neutrons

Atom Diagram, Vocab, & Scientists

Nuclear Reactions

Radioactive Decay

Alpha Particles: Nucleus releases an He (helium) atom

**weighs 4 amu less and contains two less protons

Beta Particles: Nucleus releases an electron and a neutron is converted to a proton **same mass, one more proton and one less neutron

Gamma Rays: Nucleus goes from high energy to low

**Nucleus remains the same, but a gamma ray is released

Positron: Nucleus releases a positively charged electron and converts a proton to a neutron **Same mass, one more neutron and one less proton

Electron Capture: Electron from electron cloud converts a proton to a neutron **Same mass, one more neutron and one less proton

Practice Problems

Half-Life

Half life: The interval of time required for ½ of the atomic nuclei of a radioactive sample to decay Practice Problems

Fission v Fusion

Fission: a neutron is aimed at the nucleus of a radioactive elements, causing it to split apart Fusion: nuclei collide and fuse, creating heavier elements and enormous amounts of energy Practice and More Notes

Electron Configuration

Quizlet

Energy Levels: Period Number

**The Period numbers are the ones going horizontally

S Block

● 7 energy levels

● Maximum of 2 electrons

● Last S Block formula: 1s²

P Block

● 6 energy levels

● Maximum of 6 electrons

● Last P Block formula: 2p⁶

D Block

● 4 energy levels

● Maximum of 10 electrons

● Last D Block formula: 6d¹⁰

**the D Block is always one less than the S and P Blocks, the first energy level is 3 Ex: 4s, 3d, 4p

F Block

● 2 energy levels

● Maximum of 14 electrons

● Last F Block formula: 5f¹⁴

**the first energy level is 4

Unabbreviated Electron Configuration: lists the last electron configuration of each block until it reaches the desired electron configuration (element)

**Ex: unabbreviated electron configuration of Fluorine: 1s²2s²2p⁵

Abbreviated Electron Configuration: Noble Gas Notation; Starts with the noble gas right before the element**Ex. abbreviated electron configuration of Fluorine: [He] 2s²2p⁵

d4 & d9 Exception: take away one electron from the S Block and add it to the D Block**Molybdenum: [Kr] 5s¹4d⁵

Practice with Ions, Elements, and Exceptions

Orbital Diagrams

Aufbau Principle: arrows are drawn in the sub-boxes, starting from the lowest energy level and working up Pauli Exclusion Principle: The electrons in the same orbital have opposite spins/charges Hund’s Rule: orbitals at each sub-level are half-filled before they are completely filled **Each orbital holds a maximum of 2 electrons

Drawing Orbital Diagrams Homework and Practice

Bohr Diagrams

Steps:

The atomic number represents the number of protons (which equals electrons)

The Period number of “E Levels” and the Group represents the valence electrons

Draw Element symbol

Add “E Levels” and Electrons

Drawing Bohr's Diagram Practice

Bohr’s Model Video

The Bohr Atom Video

Rutherford’s theories: Helped discover the positively charged nucleus

● Theories were wrong because charged objects that revolve in a circular motion gains acceleration ○ It is bound to liberate energy of some form

and if this continues, the energy should go

into the nucleus, creating high instability

In nature, all atoms are stable

Neils Bohr:

● only certain special orbits called discrete orbits of

electrons are allowed inside the atom

● While revolving in discrete orbits, electrons do not

radiate energy

Solar System: sun stays in the middle, planets revolve in

fixed paths

● atomic structure: the nucleus is the sun, the

electrons are the planets in orbitals

○ Obritals: energy levels

■ Energy levels indicates each shell has a defined energy level and do not liberate energy

● Knew that charged particles would give off electromagnetic radiation, so Rutherford was incorrect ○ The wavelength of the radiation varies = electromagnetic radiation spectrum

● Electrons jump between orbits

○ Did not just move back and forth, they moved like a ladder and could never be found in the middle

○ Quantized: have to be in a specific unit to exist

○ Electrons are moved when adding energy

As electrons orbit around the center, if it gains a photon, it moves up a level

As electrons orbit around the center, if it loses a photon, it moves down a level

Sir James Chadwick

● Discovered a neutral subatomic particle with the same mass as the proton

○ Neutrons (n)

Quantum Numbers

Professor Dave ExplainsClass Notes

Quantum Numbers determine the location and energy level of electrons

● Atomic orbitals- region of probability where electrons can be found ○ Holds up to 2 electrons

○ S, D, P, and F

○ More electrons = more orbitals

Principal Quantum Number (N)

● Can have any positive integer value

● Bohr Model- energy level of the electron

● Each orbital- N value

○ Higher N value = further away from the nucleus

Angular Momentum Number (L)

● Any value from 0 to (n-1)

○ Example: N=3

○ L= 0,1,2

● Describes the shape of the orbital

○ L=0, S orbitals

○ L=1, P orbitals

○ L=2, D orbitals

○ L=3, F orbitals

Magnetic Quantum Number (m_l)

● Any of value from -L to L

● Determines how many orbitals there are of a type per energy level ○ Specific orbital

Spin Quantum Number (m_s= +½ or -½)

Paul Exclusion Principle: every electron in an atom has a unique set of quantum numbers ● Any orbital can hold up to just two electrons, which have different spins

Energy Calculations

Energy Calculations Notes and Practice

Emission: releasing energy with photons

**electrons are going to a lower energy level

Absorption: absorbing energy and carrying it from a lower energy level to a higher energy level **increases energy levels

Light = energy

We can determine the amount of energy released by the light given off

E = hv

● E= Energy of the light in Joules (J)

● h= a constant, (6.626x10^-34j-s)

● v= frequency of light in Hertz (Hz)

C= λv

● C= 3.00x10⁸

● λ= wavelength

○ There are (1x10⁹ nanometers in a meter)

Periodic Trends

Understanding the Trends The Trends on the Periodic Table

Atomic Radius:

One-half the distance between the nuclei of identical atoms that are bonded together Measured in picometers

Affected by number of protons and electrons

Atomic radius decreases from left to right periods because a proton is added and the proton pulls the electron clouds in to make it smaller

Atomic radius increases as it goes down a group because you are adding an energy level further away from the nucleus

Ionic Radii

Measured in X-Rays

Ionic radii decreases from left to right periods

Ionic radii increases as it moves down a group

Anions are formed when you add valence electrons, spreading out the electron cloud Cations are formed when subtract valence electrons, shrinking the electron cloud

Ionization Energy

Energy required to remove an electron from a specific atom

Valence electrons have smaller ionization energies than inner-shell electrons

Electron shielding are the electrons added to a nucleus and the outer electrons are shielded from the nucleus Ionization energy increases from left to right periods

Ionization energy decreases as it moves down a group

Alkali metals are eager to lose an electron so that they can have the configuration of the noble gas before it `**Noble gases: very stable, hard to remove an electron energy

Atoms in which it is easy to remove the electrons have low ionization energy whereas atoms in which it is hard to remove an electron energy have high ionization energy

**Alkali metals have a low ionization energy and noble gases have high ionization energy

The electrons that are further down in groups are farther from the nucleus and it is easier to remove an electron

**lower ionization energy

Electron Affinity

Energy change that occurs when a neutral atom gains an electron

**likelihood of a neutral atom to gain an electron

Addition of an electron of a neutral atom releases energy

Electron affinity increases from left to right periods

Electron affinity decreases as it moves down a group

Halides release more energy because the addition of an electron gives it the electron configuration of a noble gas

Lower electron affinity indicates that an atom does not accept electrons as easily

Higher electron affinity indicates that an atom easily accepts electrons

First Electron Affinity vs Second Electron Affinity

**First electron affinity is the energy released when an electron is added to a neutral atom. Second electron affinity is the energy released when an electron is added to a negative ion

Second electron affinity is positive and higher than first electron affinity because it requires more energy

Depends on extent of atomic number, size of atom, stability of electron configuration **more protons = higher electron affinity

**bigger radius = lower electron affinity

Electronegativity

Measure of the ability of an atom to attract the electrons when the atom is a part of a compound Electronegativity of metals are low because they have few valence electrons

Electronegativities increase from left to right periods

Electronegativities decrease as it moves down a group

Only few noble gases will react, unknown electronegativities

Nonpolar covalent bonds: electrons shared easily

Polar covalent: no sharing of electrons

**Then forms an ionic bond

Bond Polarity

Using Electronegativity to Predict Polarity of Bonds

The property of electronegativity, which is the measure of an atom’s ability to attract electrons, can be used to predict the degree to which the bonding between atoms of two elements is ionic or covalent. The greater the electronegativity difference, the more polar the bonding is.

Polar: more time with one atom than another

Ionic bond: electrons are not shared

Polar-covalent bond: electrons are shared unequally

Non-polar covalent bond: electrons are shared equally

Metal & nonmetal bond: ionic

Non-metal & nonmetal bond: polar covalent

In polyatomic ions, both ionic and covalent bonds are present

Lewis Dot Structures

Steps for Drawing Lewis Dot Structures

Atoms

1. Write the Element Symbol

2. Find out how many electrons the atom has

3. Place one dot for each valence electron around the chemical symbol

Ex: Boron

Ions

1. Draw the lewis dot structure of neutral atom of the element

2. Think about if the atom would gain or lose electrons to fulfill the octet rule 3. Erase or add dots to represent what happened when the ion formed

4. Add the charge of the ion in the upper right corner

**cations always have 0, anions always have 8

Ex:

Bonds

Flow Chart

Drawing Lewis Dot Structures For Molecules

Bonding and VSEPR

Covalent bonds

1. Add the number of valence electrons of all the atoms in the compound

2. Arrange the atoms

a. Central atom usually has a lower electronegativity or is the first atom

3. Connect atoms with single bonds

4. Distribute electrons to fulfill the octet rule

5. Recount the electrons and bonds

Ex: C₂H₄

Ionic Bonding

Molecular Geometry

NotesVSEPR Practice Polarity Table

Valence Shell Electron Pair Repulsion (VSEPR): the VSEPR theory assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom. The five compounds show in the figure below can be used to demonstrate how the VSEPR theory can be applied to simple molecules

Two Domains around the Central Atom, Three domains around the central atom, 4 Domains around the central atom

Intermolecular Forces

NotesLabWorksheet #2Review

States of Matter

- Solids have a definite shape and volume

- Liquids have a definite volume

- Gases have no definite shape or volume

Intramolecular forces exist within a molecule

Intermolecular forces exist between molecules

London Dispersion Forces

- Occur between nonpolar molecules

- Uneven distribution of electrons = an instantaneous dipole

- Weak attractions

Dipole-Dipole Forces

- Between oppositely charged regions of polar covalent molecules

- Slightly stronger than LDFs

- Polar molecules move and rotate to maximize the attraction between oppositely charged dipoles - Strength depends on the overall polarity

Hydrogen Bonding

- Hydrogen bonds with Nitrogen, Oxygen, or Fluorine

- Strongest IMF

Phase Diagrams

NotesPractice

Heat vs Temperature

Thermal Energy- kinetic energy associated with the random motion of atoms and molecules Temperature is quantitative measurement of hot or cold

Heat (q) is the transfer of thermal energy between different bodies at different temperatures

The atoms of a solid, liquid, and gas are different

- Solid: uniform, crystalline structure, vibrate in place

- Liquid: moves around, vibrate past each other

- Gas: slides past each other, moves quickly

When atoms are heated, their kinetic energy increased and they begin to move more rapidly

Thermochemistry

PracticeCalories, calories, Joules Heat of Vaporization/FusionCombined Specific Heat: the amount of energy required to raise the temperature of 1 gram of a substance by 1 C q = mc∆T

1 Calorie = 1 Kcal

! Calorie = 1000 calories

1 calorie = 4.184 Joules

1 Calorie = 4184 Joules

Heat of Fusion: q= M∆H_fusion

Heat of Vaporization: q= M∆H_vap

Gas Laws (Again)

Part 1 ReviewPart 2Mixed PracticeTable

Boyle’s Law: (PV) = (PV)

- Indirect relationship

Charles's Law: (V/T) = (V/T)

- Direct relationship

Gay-Lussac’s Law: (P/T) = (P/T)

- Direct relationship

Combined Gas Law: (PV/T) = (PV/T)

Avogadro’s Law: (V/N) = (V/N)

- Constant pressure and temperature

- Direct relationship

Ideal Gas Law: (PV) = (NRT)

- Pressure in atm

- Volume in L

- Moles

- R: 0.08205 L*atm/Mol*K

- T in Kelvin

Dalton’s Law: P_total = P₁+P ₂…

Concentration of Solutions

PPT.Concentration CalculationsSolubility POGIL

Solvent: what is doing the dissolving

Solute: what is being dissolved

Solution: the solute and solvent combined

Grams per liter: the mass of the solute divided by the volume of the solution in liters - Concentration = (mass of solute/volume of solution)

Molarity: the concentration of a solution in moles of solute divided by liters of solution - Molarity = (moles of solute/Solution in liters)

Parts Per Million (PPM): used when low concentrations are significant, the ratio of parts of solute to one million parts solution

- Ppm = (mass of solute/mass of solution) x 10⁶

Percent Composition: the ratio of parts solute to one hundred parts solution, expressed as percent - Mass percent: (mass of solute/mass of solution) x 100%

Dilutions: take a stronger concentration of a solution, take a small amount of it, and add it to more water to create a lower concentration solution of the same substance

- (M1)(V1) = (M2)(V2)

- M1: initial molarity “stock solution”

- V1: initial volume in Liters

- M2: desired/final molarity

- V2: final volume in Liters

Solubility and Rate of Solvation

PPTPractice Worksheet

Solubility- maximum amount of solute which can be dissolved in 100g of solvent at a fixed temperature Rate of solvation- how fast does the solute dissolve in the solvent?

Saturated solution- maximum amount of solute per solvent

- Point on the curve

Unsaturated solution- less than maximum amount of solute per solvent

- Point above the curve

Saturated solution- more than maximum amount of solute per solvent

- Point below the curve

Solubility Curves

1. Find the line for the substance

2. Mass dissolved at a given temperature is on the Y-axis

3. Temperatures on are on the X-Axis

Acids and Bases

Acids and bases are pure substances

pH indicator: changes the color depending on the pH of a solution

pH scale: number scale that determines how acidic or basic a substance is (0-14)

The higher concentration of higher hydroxide ions or hydrogen ions there are, the more basic or acidic the substance is

Neutral: pH of 7 exactly

Naming Acids

- If there is no oxygen, you use the prefix :hydro” and “-ic” as a suffix

- If there is oxygen

- “-ate” because “-ic”

- “-ite” becomes “-ous”