chemistry

the diversity of matter and chemical bonding

  • Chemical bonds can be categorized into three main types: ionic, covalent, and metallic, each contributing to the unique properties of substances.

  • valence electrons are related to the group number ex Sulfur 16

  • orbitals (energy levels) are related to period number, which indicates the highest energy level occupied by electrons in an atom. For example, elements in period 2 have their valence electrons in the second energy level.


DEN- difference in electronegativity can determine the type of bond formed between two atoms, with larger differences typically leading to ionic bonds, while smaller differences result in covalent bonds.

0.0-0.3 = non polar covalent bond

0.4-1.7 = polar covalent bond

1.8 and above = ionic bond.


Metallic bonding

  • both have low electronegativity

  • share valance electrons with no chemical reaction

  • valence electrons are not held strongly by their own atoms

    → freely move between atoms

    → bonds are not easily broken if atoms are moved


Ionic bonding

  • the EN of the two atoms are quite different

  • the atom with the higher EN will remove the bonding e- from the other atom → electron transfer occurs

    → positive and negative ions are formed which electrically attract each other


electronegativity- a number that describes the relative ability of an atom to attract a pair of bonding electrons in its valence level

bonding electrons- the electrons that are involved in forming a chemical bond between two atoms

lone pair- a pair of electrons that are not involved in bonding between atoms


*Know Lewis dot, structural, and stereochemical


Molecular Polarity (diploe theory)

  • a non-polar molecule: a molecule with a symmetrical electron distribution, where all electrons are pulled equally from the center atom

  • a polar molecule: which negative charges are NOT distributed symmetrically among atoms causing a partial negative and partial positive charge on opposite sides


1) intranuclear force (bond)- bonds within the nucleus between protons and neutrons (very strong)

2) intramolecular force(bond)- bonds between atoms within the molecule or between ions within the crystal lattice(quite strong) forces within a molecule*

3) intermolecular force(bond)- bonds between molecules(weak); are electrostatic (involved positive and negative charges) forces between a molecule*

strength: inter < intra


london forces LDF = everything-weakest

diploe forces DD = polar-weak

hydrogen bonding HB = hydrogen and F,O,N- strong

LDF < DD < HB


You can not predict boiling point when:

  • one molecule has a stronger DD force and stronger LDF

  • two molecules differ significantly in shape

  • the central atom of either molecule has an incomplete octet


Surface tension:

  • a molecule within a liquid is attached by other molecules in all directions equally, but right at the surface, molecules are only attracted downwards and sideways

Capillary action:

  • due to adhesion and cohesion

    → the adhesion between water and glass is greater than the cohesion of water

    → the cohesion between mercury molecules is greater than the adhesion between the mercury and the glass


UNIT 2 -Gases


physical properties:

  • compressible

  • exert pressure on their surroundings

  • expand through diffusion

  • increase in volume with temperature, decrease in volume with pressure


[ ideal gases = obey all laws; don´t condense into liquids when cooled ]


FEMP
F: entities exert no intermolecular FORCES on one another

E: collisions between gas entities and their container are ELASTIC

M: gas entities are always MOVING in straight lines with velocities

P: gases consist of entities called POINT MASS


Temp

= the average kinetic energy of gas molecules in an area

  • the higher the Ek average, the higher the temperature

Pressure

  • atmospheric decreases as altitude increases

  • pressure = an amount of force over a unit area

Molar Volume

  • the volume one mole of gas occupies at a specific T and P

    →STP = standard temperature and pressure

    → SATP = standard ambient temperature and pressure

{ moles = volume of gas/molar volume }

n = V/Vm


Laws of Combining Volumes

  • at constant T and P, volumes of gaseous reactants are always in simple ratios of whole numbers


Avogadro´s theory

  • volume of a gas is directly proportional to the number of moles of gaseous matter present

  • at constant T and P, equal volumes of gas have an equal number of moles


Boyles law

  • in a closed system, with constant temperature and amount of gas

  • as pressure increases the volume of gas decreases

    = an inverse relationship


Charles law

  • in a closed system, with a constant pressure and amount of gas

  • as the temperature of a gas increases the volume increases

    = direct relationship


Combined gas laws

  • in a closed system, with only a constant amount of gas

  • the product of the pressure and volume of a gas is proportional to its absolute temperature

    =direct relationship, between P, V and T


ideal:

  • obeys all gas laws perfectly under all conditions

  • does NOT condense into a liquid when cooled; no attraction between particles

  • composed of point masses

  • V-T and P-T graphs yield a straight line


real:

  • does condense

  • attraction between particles at low temperatures

  • deviate from ideal gas properties at high pressure and low temperature


Unit 3- SOLUTIONS


empirical: information we gain based on what we can see, measure, record, and calculate

theoretical: information that we can interpret something; we can’t see it but we have an understanding of what must be happening


So what is a solution?

  • theoretically- the uniform mixing of entities (atoms/molecules)

  • empirically- a mixture that looks like a pure substance (salt water)


solutions = solvent + solute

→ solvent

  • where the other ¨things¨ are dissolved

  • greater quantity present

→ solute

  • the ¨thing¨ that is dissolved

  • lesser quantity present


What is an electrolyte?

  • empirical: a substance that, when dissolved in water, gives water the ability to conduct electricity

  • theoretical: a substance that when dissolved in water, forms charged entities (ions) capable of carrying/transferring electrons


Dissolving → what happens when sugar is dissolved in water?

  • polar sugar molecules are being pulled apart because of their attraction to polar water molecules

  • even things that dissociate or ionize must dissolve


Dissociating → what happens when salt is added to water?

  • polar water molecules are ¨latching¨ onto positive and negative ions in the solid and pulling them apart

  • things that ionize will dissociate but not everything that dissolves will dissociate


Ionizing → what happens when hydrogen chloride is added to water?

  • HCl was a polar molecular compound containing hydrogen instead of merely dissolving in water, the hydrogen compound forming ions


Difference:

  • dissociation occurs when ions dissociate (aka there are already ions)

  • ionization occurs when ions form ( aka from a non-ionic substance)

    → only look at acids, start with H or end in COOH

  • when trying to figure out if it is one of the three, first decide if it’s ionic or molecular


The first rule of solubility, like dissolves like

→ polar substances will be miscible or dissolve into another polar substance

*miscible- solubility but for two substances (or liquids)

→same with non-polar, but when the two meet… no mixing


→ when dealing with solid into liquid, temp increases, solubility increases

→ when dealing with a gas into a liquid, temp increases, solubility decreases


Quantitative vs Qualitative

→ quantitative

  • looks like quantities

  • deals with numbers, values, readings, math

  • later gets called analytical chemistry

→ qualitative

  • looks like quantities

  • deals with the observation of what you see (colours, reactions, burn)

  • ion colour in solution

  • flame test colours


Net ionic equations

  • it’s a reaction equation

  • the reaction between ions

  • net → we only care about the ions that change

    • if it doesn’t change we don't care = spectator ions


Concentrations refer to how much solute is dissolved into the solvent


Key Terms

  • Atom: The smallest unit of an element that retains its chemical properties.

  • Ion: A charged particle formed when an atom gains or loses electrons.

  • Isotope: Atoms of the same element with different numbers of neutrons.

  • Electronegativity: The tendency of an atom to attract electrons in a bond.

  • Molecular Shape: The 3D arrangement of atoms in a molecule determined by VSEPR theory.

  • Polarity: Uneven distribution of electron density in a bond or molecule.

Key Concepts

  • Periodic Trends: Atomic radius, ionization energy, and electronegativity.

  • Types of Bonds:

    • Ionic: Transfer of electrons between metals and nonmetals.

    • Covalent: Sharing of electrons between nonmetals.

    • Metallic: Delocalized electrons shared among metal atoms.

  • VSEPR Theory: Predicts molecular shapes based on electron pair repulsion (e.g., linear, bent, tetrahedral).

2. Forms of Matter: Gases

Key Terms

  • Kinetic Molecular Theory: Describes the behavior of gases in terms of particle motion.

  • Pressure (P): Force per unit area (measured in kPa, atm, mmHg).

  • Volume (V): Space occupied by a gas (measured in liters).

  • Temperature (T): Measure of particle kinetic energy (in Kelvin).

  • Mole (n): Amount of substance.

Key Equations

  • Boyle's Law:

  • Charles's Law:

  • Combined Gas Law:


  • Ideal Gas Law:

Key Concepts

  • STP: Standard Temperature and Pressure (0°C, 101.3 kPa).

  • Molar Volume: 22.4 L/mol at STP for ideal gases.

  • Gas behavior is ideal under low pressure and high temperature.

3. Matter as Solutions, Acids, and Bases

Key Terms

  • Solution: Homogeneous mixture of solute and solvent.

  • Solubility: Maximum amount of solute that dissolves in a solvent at a given temperature.

  • Concentration: Amount of solute per unit volume of solution (mol/L).

  • Acid: Produces ions in solution.

  • Base: Produces ions in solution.

  • Neutralization: Reaction between an acid and base to form water and a salt.

Key Equations

  • Molarity:

  • Dilution:

  • pH:

Key Concepts

  • Properties of Acids: Sour taste, conduct electricity, react with metals.

  • Properties of Bases: Bitter taste, slippery feel, conduct electricity.

  • Indicators: Substances that change color depending on pH (e.g., litmus paper).

4. Quantitative Relationships in Chemical Changes

Key Terms

  • Mole: 6.022 × 10²³ particles.

  • Molar Mass: Mass of one mole of a substance (g/mol).

  • Stoichiometry: Quantitative relationships in chemical reactions.

  • Limiting Reagent: Reactant that determines the amount of product formed.

  • Percent Yield: Measure of reaction efficiency.

Key Equations

  • Moles to Mass:

  • Stoichiometric Ratios: Use balanced equations to relate moles of reactants and products.

  • Percent Yield: ( ext{Percent Yield} = rac{ ext{Actual Yield}}{ ext{Theoretical Yield}} imes 100%

Key Concepts

  • Empirical vs. Molecular Formulas:

    • Empirical: Simplest whole-number ratio of elements.

    • Molecular: Actual number of atoms in a molecule.

  • Reaction Types: Synthesis, decomposition, single replacement, double replacement, and combustion.