Chemistry Honors⚛️👩🏽‍🔬

Unit 1: Atoms

History of the Atom

  • John Dalton- Created the theory of matter, gave the ideas that all elements (compounds) are made up of atoms. Also stated that all atoms of the same element are alike and atoms of different elements are different.

  • J.J. Thompson- He discovered the electron when experimenting with gas discharge tubes. He noticed the movement in the tube and called it the cathode rays.

  • Lord Ernest Rutherford- Conducted the gold foil experiment. He used special equipment to shoot alpha particles (positive charged particles) at the gold foil. This showed that atoms are made of a small, positive nucleus and the positive nucleus repels the positive charged particles.

  • Niels Bohr- He proposed a model of the atom to show that it is similar to the solar system. (Also called a planetary model of the atom). The electrons go around the nucleus like how planets orbit the sun.

Subatomic Particles

  • Also known as protons, neutrons, and electrons.

  • Protons have a positive charge, Neutrons have a neutral (no) charge, and electrons have a negative charge.

  • Protons and Neutrons make up the atomic mass.

  • Protons make up the atomic number. (Protons and Electrons are the same number).

  • To find how many neutrons is in an atom, subtract the atomic mass from the atomic number (Atomic mass - Atomic number)

  • Electrons have energy levels and each level can hold a certain amount of electrons. Level 1 can hold 2, level 2 can hold 8, level 3 can hold 18, level 4 can hold 32 and etc.

  • Isotopes are atoms that has the same number of protons and electrons but different number of neutrons

  • Atomic number - electrons = the charge

  • When subatomic particles have the same charge they tend to repel from each other but when it’s a different charge they move towards each other.

  • Cations are positively charger Ions and Anions are negatively charged Ions.

  • Protons and Neutrons are weighed as 1 amu

Symbols for Elements

  • “A” is another term for the atomic mass number (You can find in the top left)

  • “X” is another term to identify the element

  • “Z Number” is also the atomic number (You can find in the bottom left)

  • An isotope name can also be written as the name of the element- the atomic mass. Ex: Hydrogen-1.

Calculating Average Atomic Mass

  • There is 3 steps into calculating the Average atomic mass:

  • 1st: Convert all percentages to decimals by dividing by 100.

  • 2nd: Multiply mass of each isotope by decimal number

  • 3rd: Add everything together

  • Example: Carbon-12 is 98.93% abundant and Carbon-13 1.07% abundant. 98.93/100= 0.9893 x 12 = 11.8716. 1.07/100 = 0.0107 x 13 = 0.1391. 11.8716 + 0.1391 = 12.0107. This means that Carbon-12 is the most naturally found form.

    Fission vs Fusion

    • Nuclear fission- normally splits the nucleus into 2 nuclei, causing it to have smaller mass number aka less energy.

    • Nuclear fusion- Combines 2 nuclei to form a heavier nucleus. This causes the nucleus to gain more energy.

    • Nuclei is plural

    • Nucleus is singular

Unit 2: Periodic Table Trends

Basic Knowledge of Periodic Table

  • Periods are the rows on the periodic table

  • Groups are the columns periodic table

  • Group 1 is the Alkaline Metals. These metals are the most reactive. Their oxidation number is +1

  • Group 2 Alkaline Earth Metals. Their oxidation number is +2

  • The center of the periodic table is called transition metals. This is because the oxidation number varies.

  • Groups 3-8 are non metals.

  • Group 7 elements are called Halogens. These are the most reactive non metals. The oxidation number -1.

  • Group 8 is called Noble Gases. These are the most stable because they already have a full set of valence electrons and are octet.

    Periodic Trends

  • Atomic radii increases from top to bottom and from right to left.

  • Atomic radii has an inverse relationship with electronegativity and ionization energy.

  • Ionization energy increases from bottom to top and from left to right.

  • Electronegative also increases from bottom to top and from left to right.

  • The bigger the atom, the more reactive it is.

  • Reactivity is increases going down in metals and increases going up on non metals.

  • More protons means more attraction to electrons.

  • The greater the valence electron force means the stronger the force is.

    Unit 3: Chemical Bonds

    Different types of Bonds

    • Ionic bonds- A bond that involves a metal and a non metal. Example: NaCl. They usually fall between 2.0-4.0

    • Covalent Bonds- A bond that involves 2 non metals sharing electrons.

    • Metallic Bonds- 2 electrons move freely between metal atoms (mostly transition metals).

    • Polar covalent bonds- Electrons that are shared unequally. They usually lie between 0.4-2.0.

    • Pure (non polar) Covalent Bonds- Electrons shared equally. They fall between 0.0-0.4.

    • Polarity- Stronger positive side means it will have a weak negative side.

    Naming Molecules

  • When naming compounds, name the 1st element using the entire first name. Ex: CO- Carbon monoxide.

  • Next, name the second element using the root name + “ide” at the end. Ex: SO2- Sulfur Dioxide

  • Prefixes are used to indicate how many of each type.

    Number of Atoms

    Prefix

    1

    2

    3

    4

    5

    6

    7

    8

    9

    10

    Mono

    Di

    Tri

    Tetra

    Penta

    Hexa

    Hepta

    Octa

    Nona

    Deca

  • Examples of using the table: NF3- Nitrogen Trioxide. There are 3 fluorine atoms meaning it uses the prefix “Tri”.

  • The only time you use prefixes on the first element is when it has a subscript. Ex: As2O3- Diarsenic Trioxide. This is because Arsenic has 2 atoms and 3 Oxygen atoms.

Unit 4: Matter

States of Matter

  • Matter- Anything that has mass and takes up space.

  • Solid- A form of matter that has its own definite shape and volume. Properties: Particles that are tightly packed, expands when heated, doesn’t conform to the shape of its container, and particles move slowly. Ex: Wood, Iron, paper, and sugar.

  • Liquids- A form of matter that flows, has constant volume, and takes the shape of its container. Properties: Particles are not rigidly packed, Particles can move past each other, or flow, Liquids have constant volumes, Particles move (vibrate) more quickly. Ex: Water, Juice, and Blood.

  • The term Incompressible describes the inability of molecules in a fluid to be squeezed together (made more dense)

  • Gas- Form of matter that flows to conform to the shape of its container and fills the entire volume of its container. Properties: Particles are far away, Take the volume of their container, Take the shape of their container, Particles move (vibrate) very quickly, and gases are the only state that can be compressed easily.

    Phase Transitions

•Solid to Liquid: Melting

•Liquid to Solid: Freezing

•Liquid to Gas: Vaporization

•Gas to Liquid: Condensation

•Gas to Solid: Deposition

•Solid to Gas: Sublimation

Physical vs Chemical Properties

•Physical Properties- A property or behavior that can be observed without changing the sample’s composition. Ex: Density, Melting point, Color, Odor, Taste.

•Chemical Properties- A property or behavior that can only be observed by changing the composition of a substance. The ability of iron to form rust when combined with air and water is an example of a chemical property of iron. The inability of a substance to change into another substance is also a chemical property.

Substance vs Mixture

•Substance- Matter that has a uniform and unchanging composition. (Also known as a pure substance). Ex: Salt & Pure Water.

•Mixture- A combination of two or more pure substances. Ex: Sea Water. Sea water is a mixture because it is pure water AND dissolved salts.

•Heterogeneous Mixture: Does not blend smoothly throughout and in which the individual substances remain distinct. “Hetero”: Different. Ex: Vegetable soup, Sand and Water, Fresh squeezed Orange Juice.

•Homogeneous Mixture: Has a constant composition throughout; it always has a single phase. They blend. Ex: Tomato soup, Coffee, Chocolate ice cream

Density

  • Density: The ratio of the amount of mass in an object to the space that the object takes up (volume).

  • Density Formula: Density (D)= Mass (M)/ Volume (V)

  • The density of a substance remains the same regardless of shape, size, or mass of sample .

Viscosity

  • Viscosity: The measure of how “thick” or “sticky” a substance is and how easy it flows. Temperature may affect the viscosity of a liquid.

  • Substances with a lower viscosity tend to flow easier and be thinner. Ex: Juice & Water

  • Substances with a higher viscosity tend to flow slower and be thicker. Ex: Honey & Syrup

Unit 5: Chemical Equations

Types of Reactions

  • Synthesis: When two or more substances react and produce a single product.

  • Decomposition: A single compound that breaks down into two or more elements.

  • Combustion: When oxygen combines with a substance and release energy in the form of heat and light.

  • Single Replacement: When one element is substituted for another element in a compound.

  • Double Replacement: When two compounds exchange their components to form two new compounds.

Formulas

  • Synthesis: A+B→ AB

  • Decomposition: AB→A+B

  • Single Replacement: AB+C→AC+B

  • Double Replacement: AB+CD→AD+CB

Balancing Equations

  • The Law of Conservation of Mass: Mass cannot be created nor destroyed.

  • When balancing equations you are not creating new atoms or taking/destroying the atoms, you are simply just rearranging them.

  • Coefficients make the whole compound that same number. Ex: 6NaCl (There are 6 Sodium and Chlorine atoms now)

  • When balancing equations, if you have a subscript and a coefficient you have to multiply them.

States of Matter

  • (Aq) Aqueous

  • (L) Liquid

  • (G) Gas

  • (S) Solid

Stoichiometry

Avogadro’s Number: 6.022 × 1023 . This number is the number of units in one mole of any substance.

Moles to Particles

Moles of Given (in the question) x 6.022 × 1023 / 1 mole

Ex: How many molecules are in 14 moles of nitrogen dioxide? 14 × 6.022 × 1023 / 1 mole = 8.428 × 1024 Molecules

Particles to Moles

Particles of given (in the question) x 1 mole / 6.022 × 1023

Ex: How many moles are in 9.3 × 1024 moles of carbon dioxide? 9.3 × 1024 × 1 mole/ 6.022 × 1023 = 15.4 moles.

Mass to Moles

Mass of given (in the question) x 1 mole / Molar Mass of given