chemistry full review

Electrons shells closest to nucleus have the least amount of energy. The farther from first shell, the more energy

-ionic bonds= non metal and metal

-metallic bonds= metal only

-covalent bonds = non mental and Different Non metal

•More electron shells = larger Radius

Ions= charge due to the loss or gain of one or more electrons.

number of protons=number of electrons

Atoms= contains Protons, Neutrons and Electrons

•all atoms of an element have the same number of protons

Isotopes= diff forms of the same element, with same number of protons, and different number of neutrons

Compounds= two or more elements bind together in a fixed proportion/ratio

Mixtures= two or more substances mechanically bonded

Structural formula= identify the location of chemical bonds between the atoms of a molecule (arrangement of atoms)

Molecular formula = specifying the types and numbers of atoms present in a molecule.

Ex. C2H12O24 >> CH6O12

Chemical formulas: ex CH4

Decomposition = chemical reaction that breaks down a single compound into simpler substances or new compounds

The terms that represent two categories of compounds are ionic and molecular. 

Ionic

• the transfer of electrons between atoms, resulting in ions with opposite charges that are attracted to each other.

Molecular

• are formed when atoms share electrons through covalent bonds.

Bonding

• bonds form, energy RELEASED

• bonds break, energy ABSORBED

all atoms of an element have the same number of (PROTONS/ATOMIC NUMBER)

Solutions can be measured by

•MOLARITY

•PPM

•PBM

• a combustion reaction, where the chemical bonds are broken, releasing heat. 

Intermolecular forces: the attractive and repulsive electrostatic interactions that occur between molecules.

Ideal Gas= High Temperature, LOW PRESSURE

equal volumes of gases contain the same number of molecules. This is known as Avogadro's Law. Since both samples have the same volume of 1.0 L

A chemical reaction involves the rearrangement of atoms and molecules, resulting in the formation of new substances with different chemical properties.

Ex• O2(g) ~> O3(g)

Anode=Oxidation

Approximate mass = protons + Neutrons

The atom of fluorine forms the most polar bond with an atom of hydrogen. This is because fluorine is the most electronegative element, meaning it has a strong tendency to attract electrons in a chemical bond.

The greater the difference in electronegativity between two atoms, the more polar the bond. 

Matter

•substance(single element or combination)

Or

•mixture

Ex: air is composed of different types of gasses

The solution with the higher concentration will be a better conductor of electricity

The more ions in a solution, the lower the freezing point and the higher the boiling point

(M) = moles

Solute= what is being dissolved

Solvent= what dissolves

Solutions are HOMOGENEOUS

For solution % questions, Solute and solvent are “mass of whole” and the solute is the “mass of the part”

Water, HOH, is a neutral compound. When the volumes of the reactants combine, the moles of H'(aq) equal the moles of OH(aq) and neutralization

Entropy = disorder (Entropy is a measure of disorder or randomness in a system)

Chemical reactions occur when reactants collide with sufficient energy and the right orientation to break existing bonds and form new ones

Larger surface area = faster reaction

heat of reaction for a chemical change is (PEproducts) - (PEreactants).

the number of electrons gained must equal the number of electrons lost.

reduction is a process where a substance gains electrons, effectively decreasing its oxidation number

(LEO: Lose electrons oxidized)

(GER: gain electrons reduced)

Mass, energy and charge stay the same during chemical reactions

Electrolytic cell= electrical to chemical energy

Redox=

Oxidation is defined as the loss of electrons. When lithium is oxidized, it loses an electron to form a lithium ions

Voltaic cell= chemical to electrical

• Oxidation:
  Mg(s) → Mg²⁺(aq) + 2e⁻ (Magnesium loses electrons)
  
  
  

• Reduction:
  Ni²⁺(aq) + 2e⁻ → Ni(s) (Nickel gains electrons)

Atomic Structure

• Subatomic particles: Protons (+), neutrons (0), electrons (−)

• Atomic number = number of protons

• Mass number = protons + neutrons

• Isotopes: same number of protons, different number of neutrons

• Ions: atoms that gained/lost electrons (cations = +, anions = −)

• Electron configuration: 1s² 2s² 2p⁶, etc.

• Orbitals and energy levels

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2. The Periodic Table

• Groups = columns (same valence electrons)

• Periods = rows

• Metals, nonmetals, metalloids

• Trends:

• Atomic radius (↓ group, ← period)

• Ionization energy (↑ group, → period)

• Electronegativity (↑ group, → period)

• Reactivity trends (metals vs. nonmetals)

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3. Chemical Bonds

• Ionic bonds: metal + nonmetal (transfer electrons)

• Covalent bonds: nonmetal + nonmetal (share electrons)

• Polar vs. nonpolar

• Metallic bonds

• Lewis structures

• VSEPR theory: molecular shapes (linear, bent, trigonal planar, etc.)

• Bond polarity and molecule polarity

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4. Chemical Reactions

• Types:

• Synthesis (A + B → AB)

• Decomposition (AB → A + B)

• Single replacement (A + BC → AC + B)

• Double replacement (AB + CD → AD + CB)

• Combustion (Hydrocarbon + O₂ → CO₂ + H₂O)

• Balancing equations

• Conservation of mass

• Activity series

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5. Stoichiometry

• Mole concept: 1 mole = 6.022×10²³ particles

• Molar mass (g/mol)

• Conversions:

• g ↔ mol ↔ particles

• Mole ratios from balanced equations

• Limiting reactant & excess reactant

• Theoretical yield vs. actual yield

• % yield = (actual / theoretical) × 100

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6. Gases

• Ideal Gas Law: PV = nRT

• R = 0.0821 L·atm/mol·K

• Boyle’s Law: P₁V₁ = P₂V₂

• Charles’s Law: V₁/T₁ = V₂/T₂

• Avogadro’s Law: V ∝ n (at constant T and P)

• STP conditions: 0°C (273 K), 1 atm

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7. Solutions

• Solute/solvent

• Molarity (M) = mol/L

• Dilution: M₁V₁ = M₂V₂

• Solubility rules

• Types of solutions: saturated, unsaturated, supersaturated

• Electrolytes vs. nonelectrolytes

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8. Acids and Bases

• Properties:

• Acids: sour, pH < 7, donate H⁺

• Bases: bitter, pH > 7, donate OH⁻ or accept H⁺

• Strong vs. weak acids/bases

• pH and pOH calculations:

• pH = −log[H⁺], pOH = −log[OH⁻]

• pH + pOH = 14

• Neutralization: acid + base → salt + water

• Titration: finding concentration using volume and M

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9. Thermochemistry

• Exothermic vs. endothermic

• Heat (q) = mcΔT

• m = mass, c = specific heat, ΔT = change in temperature

• Calorimetry

• Enthalpy (ΔH): heat change at constant pressure

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10. Kinetics & Equilibrium

• Factors affecting reaction rate: temperature, concentration, surface area, catalysts

• Activation energy

• Chemical equilibrium:

• Reversible reactions

• Le Chatelier’s Principle

• Equilibrium expression: K_{eq} = \frac{[products]}{[reactants]}

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11. Redox Reactions (if covered)

• Oxidation = loss of electrons

• Reduction = gain of electrons

• Identifying oxidizing and reducing agents

• Balancing redox reactions (in acidic/basic solution)