Level 1

*know safety procedures and lab equipment

Video 1.1

• Matter: anything that has mass and takes up space

• Properties of matter:

  • Solid: definite shape, definite volume

    • This will look packed tightly together

  • Liquid: indefinite shape, definite volume

    • This will look a little looser than solids

  • Gas: indefinite shape, indefinite volume

    • This will have atoms moving around freely

• Extensive vs. intensive:

  • Extensive: amount DOES matter

    • Mass, volume, length

  • Intensive: amount DOES NOT matter

    • Color, viscosity, density, temperature

Video 1.2

• At the top you have all matter, then you ask certain questions to classify it

  • Can it be broken apart by a physical change?

    • Yes → mixture

      • Is the composition uniform/same throughout?

        • Yes → homogenous mixture/solutions (ex: salt in water)

        • No → heterogenous mixture (ex: chocolate chip cookie)

    • No → pure substance

      • Can it be separated by ordinary chemical means?

        • Yes → compound (ex: H2O)

        • No → element (ex: Na, O, H)

Video 1.3

• Physical change: identity does not change

  • State changes (ex: water to gas)

• Chemical change: identity does change

  • Signs that a chemical change has occurred:

    • Heat given off

    • PPT (precipitate) formed

    • Gas evolved

    • Color change

• Things to remember:

  • Cooking, burning, and baking something is a chemical change

  • Dissolving, crumpling, and melting something is not a chemical change

Video 1.4

• Two types of observation:

  • Quantitative: numbers evolved

  • Qualitative: no numbers, observations about characteristics

• When you solve a problem, follow this order: unit, exponent, math

• When working with scientific notation, remember “LEFTY LARGEY” and “RIGHTY REDUCEY'“

  • Lefty largey: 300,000 → 3 • 10^5

  • Righty reducey: 0.00089 → 8.9 • 10^-4

• MATH WITH EXPONENTS:

  • Multiplication: ADD the exponents together

  • Division: SUBTRACT the exponents

  • Addition and subtraction: make the exponents the SAME

Video 1.5

• Significant figures are used to indicate how accurate your measurements are

• We always record a measurement with all the numbers we know with certainty plus one uncertain guess

  • Always go beyond one marking (if the smallest markings are by ten, your guess will be in the ones place)

• Rules for counting significant figures:

  • All non-zero digits are significant

    • 1,781 = 4 sig figs

  • Leading zeros are never significant

    • 0.056 = 2 sig figs

  • Captive zeros are always significant

    • 1,002 = 4 sig figs

  • Trailing zeros are significant if a decimal is present

    • 300.0 = 4 sig figs

  • Exact numbers have infinite sig figs (if the measurement is not EXACT, this will mess up your sig figs so be careful)

    • 1,000 mL = 1 L

Video 1.6

• When doing multiplication and division for sig figs, remember the word “LEAST”

• When doing addition and subtraction for sig figs, remember to count up the columns and make the scientific notation exponents the same

• Remember units, exponents, sig fig

• Remember how to do multiplication, division, and subtraction/addition with scientific notation

• When doing multiplication and division, you will also add or subtract the units as well

Video 1.7

• Dimensions:

  • Length: meter, feet

  • Volume: liter, gallon (this will have a unit raised to the third power)

  • Mass: kilogram, pound

  • Density: kilograms per liter, pounds per cubic foot (this will have a mass over something cubed)

  • Temperature: Celsius, Kelvin, Fahrenheit

  • Time: seconds, hour

  • Pressure: KPa, pound per inch squared

  • Energy: joule, calorie

  • Quantity: mole, mol

• Celsius to Kelvin: C + 273.15

• Kelvin to Celsius: K - 273.15

• Calculating percent error:

  • Absolute value of experimental value - actual value, divided by actual value, multiplied by 100

  • You will have a NEGATIVE percent error, if your measured value was LESS than the actual value

  • You will have a POSITIVE percent error, if your measured value was MORE than the actual value

Video 1.7b

• Prefix conversions:

  • Exa, E, 10^18

  • Peta, P, 10^15

  • Tera, T, 10^12

  • Giga, G, 10^9

  • Mega, M, 10^6

  • Kilo, k, 10³

  • Hecto, h, 10²

  • Deka, da, 10^1

  • Deci, d, 10^-1

  • Centi, c, 10^-2

  • Milli, m, 10^-3

  • Micro, u, 10^-6

  • Nano, n, 10^-9

  • Pico, p, 10^-12

  • Femto, f, 10^-15

  • Atto, a, 10^-18

• When you are converting with prefixes, just subtract the scientific notation exponents

  • ex: 15.6 T → u (10^12 - 10^-6 = 10^18) = 1.56 × 10^19

Video 1.8-1.9

• When converting quantities, follow this order:

  • Write given amount as fraction

  • If one dimension, put fraction over one

  • Choose conversion factors from table

  • Write conversion factors to cancel unit

  • Cancel units to get desired units

  • All numbers on top, you multiply; all numbers on bottom, you divide

• Sig figs will usually be the original number you worked with

• When dealing with multiple dimensions, follow this order:

  • Do one dimension at a time

  • When units on bottom, cancel up

  • When dealing with square or cubic units, you must write each twice or three times

Level 2

Video 2.1

• Dalton

  • Matter is composed of small particles (atoms)

  • Atoms are identical if they are the same element; different atoms have different properties

  • Atoms cannot be created or destroyed

  • Atoms combine in whole number ratios

  • In chemical reactions, atoms combine/rearrange/separate

• Thomson

  • Known for his cathode ray tube experiment

  • He used a piece of metal, which was made out of atoms, and put it in the cathode ray tube

  • When the ray shot out, he noticed that it bent away from one of the metal plates, which was negatively charged

  • Showed that atoms were NOT solid balls, but actually spherical things that have CHARGES

  • Known for his plum pudding model, where you have electrons and protons in an unorganized lump

  • PROOF OF ELECTRONS

• Rutherford

  • Known for the Rutherford experiment

  • Used gold foil and shot alpha particles (positives) at it, while also putting a fluorescent screen to be able to see the atoms

  • Noticed there would be a random hit to the side every few thousand times

  • Positives went right through the gold foil, implying that the positives and negatives could NOT be randomly dispersed

  • Rutherford model where protons are grouped in the middle (nucleus) with electrons dispersed

  • PROOF OF NUCLEUS

• Chadwick

  • Came up with dead weight

  • Figured out the masses of atoms because protons and electrons did not work because it would not match up

  • PROOF OF NEUTRONS

• Bohr Model

  • Works because atoms jump up and fall down when given energy, which can be seen in the form of light; this makes lines that represent and energy fall down

  • Seven energy level: inner shell = 2, second shell = 8

Video 2.2

• If you blew up each atom to the size of a blueberry and filled a grapefruit, the grapefruit would have to be the same size as the earth. If you filled the earth with blueberries, you would have the same number of, for instance Nitrogen atoms, as a grapefruit.

• The nucleus in an atom is insanely small. If you blew up the atom as a whole into the size of a blueberry, you would still not be able to see the nucleus. If you blew up the blueberry/atom into the size of a two-story house, then the nucleus would be barely visible. You would have to blow up the blueberry/atom into the size of a football stadium to see the nucleus as the size of a marble.

• The nucleus of an atom is incredibly dense. To give a comparison, you would actually need to pack 6.2 billion cars into a 1×1×1 ft box to model the same density as an atom’s nucleus

• Atoms are made of protons, electrons, and neutrons

• There is empty space between nucleus and electrons

• Electrons weigh basically nothing

Video 2.3

• Atomic number: number of protons and number of electrons in a NEUTRAL atom

  • Number of protons determine the IDENTITY

• Mass number: Neutrons and protons, entirety of the mass is essentially in the nucleus

• Average atomic mass: certain atoms of certain elements have different masses, so it is averaged out sometimes (% • mass number)

• Isotope: same number of protons, but different number of electrons

• Sig fig for periodic table = 2 decimal places

• Atoms are electrically neutral because they have the same number of positively charged protons and negatively charged electrons

Video 2.4

• Two main categories of elements on the periodic table:

  • Non-metals

  • Metals

• Properties of metals:

  • Shiny/luster, malleable, conductors, ductility

• Properties of non-metals:

  • Dull, not malleable, brittle/powdery, cannot bend

• Metalloids: elements that have properties that are in-between metals and non-metals

• Where to find metalloids: jumps along the staircase

• Two elements that are liquid at room temperature:

  • Bromine

  • Mercury

Video 2.5

• Valence level of electrons: outermost energy level, most important level in the atom because it interacts with other atoms (transferred and shared)

• Atoms always try to get an OCTET (8 electrons in outershell)

• Atoms that lose an electron gets a positive charge because electrons are negatively charged. The positive charge comes from the fact that the atom has more protons than electrons

• Atoms that gain an electron gets a negative charge because there are more electrons, which are negatively charged, than protons, which are positively charged.

• Remember these examples:

  • Bromine in its natural state has seven electrons, but after gaining one valence electrons, becomes Br^-

  • Calcium in its natural state has two valence electrons. After losing these two electrons, it becomes Ca²+

  • Think opposite of charge when it asks you how much electrons it will gain or lose

Video 2.6

• Chemical bond: a connection between two atoms

• Three types of bonds:

  • Ionic: metal with non-metal (ex: MgCl2)

  • Covalent: non-metal with non-metal (CO2)

  • Metallic: metal with metal (Na)

• Ionic bonds are caused by stealing valence electrons. One atom transfers electrons to another, resulting in a positively charged ion and a negatively charged ion that are attracted to each other

• Covalent bonds are caused by sharing valence electrons. Both atoms are attracted to the same pair of electrons, creating a stable connection between them

• Electronegativity: how able an element is to gain an electron

• Diatomic atoms: BrINClHOF

Video 2.7a-2.7b:

• Cation: an atom or a group of atoms with a positive charge

  • Metals

• Anion: an atom or a group of atoms with a negative charge

  • Non-metals

• Writing formulas of ionic compounds:

  • Write the symbol of the cation with its charge superscripted

  • Write the symbol of the anion with its charge superscripted

  • Find the ratio of cation to anion that cancels all charges

• Examples:

  • Magnesium chloride: MgCl2

  • Aluminum oxide: Al2O3

  • Lead (II) chloride: PbCl2

• Naming ionic compounds:

  • Name the cation

  • Use a roman numeral if the element is a transition metal

  • Name the anion

• Examples:

  • CaF2: calcium fluoride

  • Al2S3: aluminum sulfide

  • CoBr2: Cobalt (II) bromide

• Remember that positives always go first and that you can ALWAYS find the charge of the anion

Video 2.8a-2.8b:

• Polyatomic ion: a group of atoms with a charge

• Writing formulas of ionic compounds with polyatomic ions examples:

  • Magnesium nitrate: Mg(NO3)2

  • Ammonium carbonate: (NH4)2CO3

  • Lead (IV) hydroxide: Pb(OH)4

• Naming ionic compounds with polyatomic ions examples:

  • CaSO4: calcium sulfate

  • Cu2CO3: copper (I) carbonate

  • (NH4)3PO3: ammonium phosphite

Video 2.9:

• Covalent/molecular compound: two or more non-metals sharing electrons together

• USE THE GREEK SYSTEM

  • Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

• Naming rules:

  • The first element (if only one), write the name of the element

• Examples:

  • CO2: carbon dioxide

  • N2O5: dinitrogen pentoxide

  • CO: carbon monoxide

  • SiO4: silicon tetroxide

Video 2.10:

• Acids always have HYDROGEN and NO CHARGE in them

• -ide → hydro_ic

  • ex: Hydrogen + bromide → hydrobromic acid (HBr)

• -ate → _ic

  • ex: Hydrogen + nitrate → nitric acid (HNO3)

• -ite → _ous

  • ex: Hydrogen + bromite → bromous acid (HBrO)

Level 3

*Some videos are skipped because you do not need to know them

Video 3.3

• Rules for balancing equations:

  • Do not mess with a compound’s structure by changing the subscripts

  • Only change COEFFICIENTS to balance

  • Start with the atoms that are in only one compound on each side

  • Treat polyatomic ions as a single unit

  • Leave oxygen and hydrogen until the very end

    • CHO2 rule (do carbon, then hydrogen, then oxygen, and if you have an ODD number of oxygens, multiply everything by 2)

  • DOUBLE CHECK

Video 3.4

• Synthesis reaction:

  • A + X → AX

  • Taking 2 individual things and putting them together to form something else

  • ex: 2Mg + O2 → 2MgO

• Decomposition reaction:

  • AX → A + X

  • Opposite of a synthesis

  • ex: H2O → 2H2 + O2

• Single replacement reaction:

  • A + BX → AX + B

  • MUST USE ACTIVITY SERIES

  • ex: Al + Cu(NO3)2 → Al(NO3)3 + Cu

• Double replacement reaction:

  • AX + BY → AY + BX

  • USE SOLUBILITY RULES

• Combustion reaction:

  • CHO2 rule

  • Reaction with oxygen

    • Must be an alkane (anything with carbon, hydrogen, and/or oxygen) (take carbon, multiply it by 2, then add 2 to find hydrogen)

      • Methane: CH4

      • Ethane: C2H6

      • Propane: C3H8

      • Butane: C4H10

      • Pentane: C5H12

      • Hexane: C6H14

      • Heptane: C7H16

      • Octane: C8H18

      • Nonane: C9H20

      • Decane: C10H22

  • Products will ALWAYS be CO2 + H2O

Extra notes to know for level 3:

• When doing solubility rules for double replacement reactions, INSOLUBLE is what you want

  • When writing NET IONIC EQUATIONS, you NEED to put the charge

• Know how to name and balance all types of reactions

• WEIRD DECOMPOSITION REACTIONS (where M represents a metal):

  • MCO3 → MO + CO2 (carbonate)

  • MOH → MO + H2O (hydroxide)

  • MClO3 → MCl + O2 (chlorate)

• WEIRD DECOMPOSITION REACTIONS WITH ACIDS:

  • H2SO3 → SO2 + H2O

  • H2CO3 → CO2 + H2O

• WEIRD SYNTHESIS REACTIONS (where M represents a metal):

  • MO + CO2 → MCO3

  • MO + H2O → MOH

  • MCl + O2 → MClO3

Extra information that will be useful:

• Structure normally associated with a covalent network solid:

  • Atoms covalently bonded in a lattice structure

    • Atoms in a covalent network solid want to maximize the number of strong covalent bonds they form. Arranging themselves in a lattice lets them do this efficiently, creating a stable and strong material.

• Ionic (molecular) compounds do not conduct electricity because…

  • Electrons are not free to move in an ionic crystalline lattice

    • Ionic bonds are formed by the electrostatic attraction between positively charged ions (cations) and negatively charged ions (anions)

    • A lattice structure arranges these ions so that every cation is surrounded by as many anions as possible, and vice versa. This maximizes the attractive forces, making the structure more stable.

• Structure normally associated with a covalent bonds:

  • A simple molecule

    • A simple molecule is a small group of atoms bonded together by covalent bonds, where atoms share electrons to achieve a stable outer electron shell

    • These molecules have a distinct, finite number of atoms (unlike covalent network solids, which have an extended structure)

• Sodium chloride not conduct electricity in its solid form because…

  • Ions in its structure are in fixed positions

    • In the solid state, sodium chloride forms a crystalline lattice structure, where sodium ions (Na⁺) and chloride ions (Cl⁻) are arranged in a repeating pattern. These ions are held in fixed positions by the strong ionic bonds between them

    • For electricity to flow, there must be charged particles (ions or electrons) that are free to move and carry the electrical current

    • In solid NaCl, the ions are locked in place within the lattice and cannot move around. Therefore, there are no free-moving ions available to conduct electricity