Chemistry - Sem1 Final

Democritus

440 BCE proposed that everything in the world was made out of particles: Atoms (indivisible)

John Dalton

\`1808: Atomic Theory (Accepted)


1. various compounds are combinations of atoms of differing elements
2. elements are of differeing size and mass
3. elements cannot be created nor destroyed
4. elements join together to create many different things

JJ Thomson

1897 discovers the electron; creates the “Plum Pudding Model”

Ernest Rutherford

“Father of the nuclear age”: though shooting alpha particles through gold foil, conclude that atomse had a lot of empty space with its mass at the center.

Niels Bohr

Electrons orbit the nuclues at fixes energies and distances. The can jumpdetwee the levels but cannot live in the space between levels. An atom is a “planetary model,” with the electrons acting as planets, with the abaily of a quantum leap

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Quantum Leap

Niels Bhor, 1897

* “Quantum Leap”: As electron jump between “orbits” they lose a descrete amount of energy. This is known as a Quantum of energy. The more energy an electron has, the further away it is from the nucleus

Werner Heisenberg

Impossible to determine the position ans speed of electrons

* Quantum model

Atomic Number

Displays number of protons in an element

Isotopes

Atoms of a given element differing in number of neutrons

Mass Number

Number of protons + number of neutrons

Ions

Atoms that have gained or lost electrons

Cation

Positively charged ion

Anion

Negatively charged ion

Ionic Compound

Elements are held togeht by a difference of charge (one element gains electrons from the other element)

Covalent Compound

Bonding by sharing electrons

Atomic symbol

Matter

Anything that has mass and takes up space

Chemistry

The study of matter; almost everything outside of energy

Physical Proprty

Observed with the 5 senses

Chemical Property

Observation needing more than just the 5 senses

Physical Change

Change without changing the identity of a substance. Doesn’t change ts stucture/properties

* phase changes, breaking, tearing, dissolving

Chemical Change

Matter changes its identity. Change in composition of a substance in a reactions.

* Release of energy (light/heat), color change, forms gas/precipitate, not easily reversed

Mixtures

Combination of matter with variable composition. Separated physically

* Distillation (sugar + water)
* Homogeneous: Uniform composition; equal throuhgout
* Heterogeneous: Depending on location within mixture

Pure Substance

Not separated physically, if at all; same ratio of components but differing identities

* Elements: Cannot be separated chemically
* Compound: 1 or more elements combined; separated chemically

Metric System

1 Meter:

* 1,000 milimeters
* 100 centimeters
* 10 decimeters
* .10 dekameters
* .001 hectometers
* .0001 kilometers
* Mili: 1/1000
* Centi: 1/100
* Deci: 1/10
* Base (meter, gram, liter): 1
* Deka: 10
* Hecto: 100
* Kilo: 1000

Factor-Label System

Using units to solve measurement problems

Significant figures

Sig Figs; Numbers you can certify/guarantee

* Yes: non-zero numbers, Captive zeros, trailing zeros with a decimal
* No: Leading zeros, trailing zeros without a decimal
* Add/subtract: round to lost decimal place
* Multiply/divide: Round to lowest sig figs

Accuracy

How close a measurement is to the accepted value

* Universal

Precision

How close measurements of the smae item are to eachother

* Subjective

Density

Mass / Volume (M/V)

Naming Molecular Formulas

Nonmetals (4A - 7A)


1. Determine the names of the symbols
2. Determine the number of atoms of each element present
3. Write prefix corresponding with the quantity of each atom
4. End last element by changing the ending to -ide

Writing Molecular Formulas

Nonmetals (4A - 7A)


1. Write sumbol for each element
2. Determin the subscribt from the prefix of each element
3. CANNOT REDUCE

Naming Ionic I Formulas

Metals with only 1 charge (1A - 3A)


1. Name the metal and nonmenal
2. End nonmetal with -ide
3. NO PREFIXES OR NUMERALS

Writing Ionic I Formulas

Metals with only 1 charge (1A - 3A)


1. Write symbol for metal and nonmetal
2. Find both charges from the Periodic Table
3. Equalize with Criss-Cross
4. CAN REDUCE

Naming Ionic Type II Formulas

Transition Metals w/ 1+ Charges


1. Name Metal and Nonmetal
2. Use __**Heeren’s Law**__ to find the charge


1. Multiply the charge of nonmetal by its subscript Add/Subtract it from rigth + left, divide right by subscript of metal
2. Subscript Metal x Charge Metal) + (Subscript Nonmetal x Charge Nonmetal) = 0
3. Use Roman Numerals to represent charge
4. End name of nonmetal in -ide

Writing Ionic Type II Formulas

Transition Metal w/ 1+ Charges


1. Write Symbol for metal and nonmetal
2. Determine charge from Periodic Table/Numeral
3. Criss-Cross charges
4. CAN REDUCE

Combination Reaction

Combo; 2 Reactants to 1 Product


1. A + B = AB


1. Always check charges and balance

Decomposition

Decomp; 1 Reactant to 2 Products


1. AB = A + B


1. Check charges and balance

Combustion

C.C.; Burning, needs fuel source


1. C_H_ + O2 = CO2 + H20


1. CO4 = Methane

Single Replacement

SR; Element + Compound to New Element + New Compound


1. AB + C = AC + B


1. ALWAYS check activity series; only a higher-ranked element can replace a lower-ranked metal

Double Replacement

DR; 2 Ionic Compounds to 2 New Compounds


1. AB + CD = AD + CB


1. Usually forms a precipitat, gas (bubbling), or water

Mole

SI (Metric) Standard unit to measure the amount of a substance

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1 mole is equivalent to:


1. Atomic mass of an element
2. 6.02e23 atoms (∞) - __*Avogadro’s Number*__
3. 22.4L (∞)

Dimentional Analysis

Relationship between quantities based on fundamental quanlity


1. State Given
2. Multiply by conversion factor (atomic mass, atoms, liter)
3. Arrange Step 2 so desire units are on top
4. Solve and round to sig figs

Molar Mass

The molar mass of a total compound


1. Find number of atoms of each element
2. Find mass of each individual atom


1. From Periodic Table
3. Find total mass of the element available
4. Find total mass of compound


1. Add up step three of each element

Percent Composition

Comparison of all elements in a compound : compound mass


1. (Mass of element) / (Mass of compound) x 100

Hydrates: Finding anhydrate

Hydrate: Substance w/ water as an additional component


1. Find molar mass of each reactant
2. Add up molar masses to find hydrate mass
3. Divide molar mass of individual reactant by hydrate mass
4. Divide by 100, multiply by quantity of anhydrate

Hydrates: Coefficient of Hydration

Mole Ration


1. Find mass of both reactants in substance
2. Convert masses of both into moles
3. Divide both moles quantities by the smallest mole quatitity
4. Quotients are the equtation coeffients


1. Quotient that is >1 should be your coefficient of hydration

Empirical Formula

A formula giving the __proportions__ of the elements present in a compound but not the actual numbers or arrangement of __atoms__.


1. If given percentages, assume that you have a 100g sample
2. Convert percentages into grams
3. Convert mass into moles
4. Divide quantities by the smallest mole quantity
5. Results are the equation subscripts 


1. Is the basic, lowest ratio that you can have for a substance

Molecular Formula

The real compound that still shares the same ratio as its empirical formula


1. A multiple of the empirical formula

Stoichiometry

Relating the Amount of one substance to another


1. Balance Chemical equation
2. Convert given into moles
3. Convery step 2 into moles of desired substance


1. Use mol/mol ration using moles from equation
4. Convert step 3 into desired unit

Percent Composition

The percentage your answer is from the theoretical result


1. (Actual yield) / (Theoretical yield) x 100


1. Normally

Limiting Reactant

You are limitd by whatever runs out first


1. Balance given equation
2. Use Stoichiometry to determine how much of your porduct each substance can make
3. Whichever theoretical yield is lower is your limiting reactant

Properties of gases

1. Compressible: Able to be compressed/moveable
2. Expanding: Will take up full volume of structure
3. Shape: Takes shape of its container
4. Low Density: Many float (Helium)
5. Homogeneous: Equal distribution within a mixture
6. Similarity: Properties of all gases are similar

Kinetic Molecular Theory

Takes idea of what we can see in order to explain what we cannot see


1. Gas particles are so small that their volume is considered to be 0L
2. Gas particles are in constant motion.


1. Collisions with wall creates pressure
3. Gas particles do not attrct or repel eachother
4. Average kinetic energy is proportional to temperature


1. Cold Temperature: Less energy
2. Warm Temperature: More Energy

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Boyle’s Law

Pressure and Volume are inversely related when temperature is constant


1. P1V1 = P2V2

Charle’s Law

If temperature of a gas increases, its volume increases IF pressure reamins the same

V1 / T1 = V2 / T2

Temperature

Measure of Average Kinetic Energy of Particles


1. Kelvin: Starts at absolute zero


1. Absolute zero: All particles stop moving
2. K = °C + 273
3. Used in ALL gas-related problems

Combined Gas Law

Combination of Boyle’s Law and Charle’s Law; When everything changes


1. Anything constant will cancel

STP

Standard Temperature and Pressure


1. P = 1 atmosphere or its equivalents
2. T = 273 K (0 °C)

Dalton’s Law of Partical Pressures

The pressure of each gas is proportional to the moles available


1. “x” : Mole fraction - Mole gas 1 / total moles
2. P(gas1) = (X - gas 1)(P - total)
3. P -Total = Sum of pressures of each gas

Avogadros Law

Volume is directly proportional to the number of gas molecules if Pressure and Temperature are constant


1. From Gay-Lussac: Gases comine in volumes that have smmple whole number ratios; pressure and temperature are directly proportional

Ideal Gas Law

Ideal Gas: Occupy ngligible space and have no interactions; obeys all gas laws.


1. 1 situation; NOT CHANGING
2. PV = nRT


1. P = Pressure
2. V = Volume
3. n = Moles of gass
4. R = 0.08206 (when P = atm); 8.314 (when P = kPa); 62.36 (when P = Torr)
5. T = Temperature

Ideal Gas Law: Molar Mass

Ideal Gas Law: Density

calorie

cal: Heat required to raise 1g water by 1 °C

* 1 cal = 4.184 J
* 1 cal = 0.0001 kcal

Calorie/Kilocalorie

Cal/kcal; Heat required to raise 1,000g water by 1 °C

* 1 Cal = 4,184 J
* 1 Cal = 1,000 cal

Joule

J; Standard SI unit for heat/energy

* 4.184 J = 1 cal
* 4,184 J = 1 Cal/ 1 kcal

Specific Heart

* Specific Heat Capacity = C
* Quantity of heat (J, cal, Cal/kcal) absorbed per unit mass (g, kg) when temperature increases / decreases by 1 °C
* Reason why some materials hold heat longer / change heat faster
* Higher C = More energy to increase temperature

Change in Heat - ΔH

Equation to find change in heat

* m = Mass
* C = Specific Heat
* ΔT = Change in temperature (final temp - first temp)

Equilibrium

mCΔT = -mCΔt

Occurs when liquid is added to a calorimeter

Calorimetry

* Calorimeter: Device to measure heat of a given object/reaction
* Heat lost = -Heat gained
* Heat lost by an object is gained by its surroundings
* Detects calories in food

0th Law

2 thermodynamic systems are in equilibrium with another

* |A|B|C|
* (A + C) / 2 = B
* Underlying principle in temperature found in other laws
* found after Laws 1-3

1st Law

As heat increases, so does the amount of potential work

* W = P(ΔV); P = Pa; V = m3

Carnot Efficiency

Efficiency of a system is based off of the difference between hot and cold temperatures

* E = \[(T-hot ) - (T-cold)\]/T-hot x 100

2nd Law

Heat flows from Hot to cold

* Can flow in opposite direct but needs a lot of help

Entropy

Amount of energy given off with every change in temperature

* Measure of the disorder within a system (gs has more entropy than liquid)

3rd Law

As temperature approaches absolute zero (0k) entropy is minimized and all functions stop

* Cannot do Charle’s Law at 0 K

State Change

* Freezing: Process of taking a liquid to a solid (0 °C)
* Melting: Process of taking a solid to a liquid (0 °C)
* Evaporation: Process of taking a liquid to a gas (100 °C)
* Condensation: Process of taking a gas to a liquid (100 °C)
* Sublimation: Process of taking a solid to a gas

Heating / Cooling Curve

Heating: Graph of phase changes thorugh Increasing Temperature

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Cooling: Graph of phase changes through Decreasing Temperatures

Heat of Vaporation and Fusion

Vaporation: Amount of energy to boil a liquid

* Hvap
* Water = 2,260 J/g
* 100 °C

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Fusion: Amount of energy to freeze a liquid

* Hfusion (Hf)
* Water: 333 J/g
* 0 °C

Heat Change with Phase Change

During a phase change, you have to account for any phase changes that occur

* ΔH = (mCΔT) + H(v or f)
* Can have multipl mCΔT’s and H(v or f)’s depending on how many phase changes occur

Enthalpy

The sum of internal energy and product of pressure and volume

* 1st Law
* Directioni of heat transfer in a reaction; ΔH Products - ΔH Reactants
* Endothermic: Heat absorbed (+)
* Exothermic: Heat release (-)
* Use enthalpy to fin the ΔH of variables

Hess’s Law

Enthalpu is the function of the state of matter

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If desired reaction is multiple steps, show the enthalpy for each step

Solving Hess’s Law equations

* Undesired variables: Put on opposite sides
* Flip: Products to reactants and vice versa; change sign of given ethaply
* Desired Variables: On desired side with desired quantity
* Multiply the whole equation by factor to get desired quantity; given enthaply is multiplied by same factor
* Add up your ending enthalpies to find final enthalpu

The History of Quantum Mechanics

* Democritus: 400 BC - Proposed the “atom”
* Dalton: 1803 - Atomic Theory
* Small Particles
* Indivisible
* Unique
* Comine and Separate
* Golstein: 1886 - “Anoderays” lead to the discovery of the Proton
* Thomson: 1897 - “Cathoderays” lead to the discovery of the electrom
* Plum Pudding Model
* Rutherford: 1911 - Discovers te nucleus during the Gold Foil Experiement

Particles and Waves

Particles: Electons are assumed to be particles

* Particles have mass and occupy space
* Very small mass (9.11e-31 kg)

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Waves have no mass but carry energy

* If electons are “massless” can they carry energy?

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Much knowledge of the electron comes from the study of light

Light

* Colors: Absorbed and Reflected
* Fast, no mass, constant motion
* Electromagnetic Spectrum: R. O. Y. G. B. I. V.

Photon

Quantum of light carrying energy proportional to radiation frequency; NO MASS

Frequency

* 𝜈: (nu); represents frequency
* Number of waves travelling every second
* Hertz: Measurement of Frequency
* 1 Hz = 1 wave per second
* kHx (kilohertz) and MHz (megahertz)

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INVERSE OF WAVELENGTH

Wavelength

λ: (lambda); symbol of wavelength

* Distance between 2 consecutive crest/troughs of a wave

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Monometers: Measurement of a wavelength

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INVERSE OF FREQUENCY

Speed of light

* c: Speed of Light
* All light travels at same speed
* 3.00e8 m/s
* Sp = λ • 𝜈

Wave - Particle of Light

Different color = different wavelength = different frequency = different energy

* Energ of light travels in “packets” or “quanta” which cna be measure by function of frequency

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E = h𝜈

* E = Energy
* H = Planck’s Constant: 6.626e035 j/sc

Color: Photon Energy

Emission Spectra

* A spectroscope separates the light into bands (lines) of specific wavelengths
* light given off = emission spectrum
* light energy gained = absorption spectrum
* Each element has a unique spectrum which can be used to identify it

Excitement

* When atom gains energy, electron “jumps” to a higher energy level
* Also called quantum level
* As the  electron “drops” to a lower energy level, the energy is released in the form of light

Atomic Spectra

* Atoms may gain extra energy - become excited - and they release that energy in the form of light


* EX: Luminol, Glow Sticks

Quantum Theory

* Atoms emit or absorb only radiation with certain specific frequencies.
* Energy is directly proportional to frequency
* So only certain energy states/levels of the electrons are possible 
* “Quantized”

Electromagnetic Radiation and Color

* Compounds absorb light when the difference in energy levels corresponds to a wavelength that is visible to the eye.
* When some ionic compounds are heated, the heat of the flame provides enough energy to excite an electron to a higher level.
* Then the light is emitted when the electron falls back.

Bohr Model

Electron “orbi” the nucleus like planets orbit the sun

* Held together by force of attraction between opposite forces

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Energy related to the distance from the nucleus (close = higher attraction)

Bohr Model Diagram

1. 2
2. 8
3. 8
4. 18
5. 18
6. 32
7. 32