CHEM 131 Midterm 1 Notes

Dalton’s Theory of the Atom

Elements are made up of tiny, indestructible particles called atoms.

Atoms of a given element

All atoms of a given element have the same mass and distinctive properties.

Combining Atoms

Atoms combine in whole ratios to form compounds.

Proton Mass (kg)

Mass of a proton is 1.67262 × 10^-27 kg.

Proton Mass (amu)

Mass of a proton is 1.00727 amu.

Proton Charge

Relative charge of a proton is 1+.

Electron Mass (kg)

Mass of an electron is 0.00091 × 10^-27 kg.

Electron Mass (amu)

Mass of an electron is 0.00055 amu.

Electron Charge

Relative charge of an electron is 1-.

Neutron Mass (kg)

Mass of a neutron is 1.67493 × 10^-27 kg.

Neutron Mass (amu)

Mass of a neutron is 1.00866 amu.

Neutron Charge

Relative charge of a neutron is 0.

Isotopes

Elements with the same number of protons but different neutrons are called isotopes.

Atomic Number (Z)

The atomic number indicates the number of protons in an element.

Element Symbol

The symbol used to represent an element (e.g., C for Carbon).

Atomic Weight

The atomic weight is the weighted average mass of an element's isotopes.

Ions

Ions are charged atoms formed when atoms gain or lose electrons.

Neutral Atom

In a neutral atom, the number of electrons equals the number of protons.

Cations

Positively charged ions are called cations.

Anions

Negatively charged ions are called anions.

The Plum Pudding Model

Atoms contain negatively charged electrons within a positively charged space.

Nuclear Model of the Atom

The nucleus contains most of the atom's mass and positive charge.

Bohr Model

Electrons travel in fixed orbits around the nucleus corresponding to specific energy levels.

Quantum Mechanical Model

Electrons exist in probability domains rather than fixed orbits.

Wave-Particle Duality

Electrons exhibit both wave-like and particle-like behavior.

Schrodinger’s Equation

An equation used to calculate the probability of finding an electron in a specific region.

Principal Quantum Number (n)

The principal quantum number defines the energy level of an electron.

Angular Momentum Quantum Number (l)

The angular momentum quantum number indicates the shape of the orbital.

Magnetic Quantum Number (ml)

Indicates the orientation of an orbital within a given sublevel.

Pauli Exclusion Principle

No two electrons can have the same four quantum numbers.

Coulomb’s Law

Like charges repel and unlike charges attract; energy lowers when charges separate.

Effective Nuclear Charge (Z-effective)

The net positive charge experienced by an electron in the atom.

Shielding Effect

Inner electrons shield outer electrons from the full nuclear charge.

Penetration Effect

The ability of an electron to get closer to the nucleus, affecting its energy.

Aufbau Principle

Electrons fill orbitals from lowest energy to highest.

Hund’s Rule

Electrons occupy empty orbitals before pairing up in the same orbital.

Valence Electrons

Electrons in the outermost shell that determine an element's chemical properties.

Core Electrons

Electrons in the inner shells that do not participate in chemical bonding.

Ion Formation

Atoms become ions by gaining or losing electrons to achieve full valence shells.

Periodic Law

Elements with similar properties recur at regular intervals and fall into columns.

Alkaline Earth Metals

Elements in column 2A of the periodic table.

Alkali Metals

Elements in column 1A of the periodic table (excluding hydrogen).

Halogens

Elements in column 17 of the periodic table, they tend to gain one electron.

Noble Gases

Elements in column 18 of the periodic table are unreactive due to full valence shells.

Transition Metals

Elements located in columns 3-12 of the periodic table.

Electron Configuration Exceptions

Certain elements exhibit unexpected electron configurations for stability.

Ionization Energy (IE)

The energy required to remove an electron from an atom in its gaseous state.

Electron Affinity (EA)

The energy change associated with the gaining of an electron by an atom.

Ionic Bonding

A bond formed through the electrostatic attraction between oppositely charged ions.

Covalent Bonds

Bonds formed by the sharing of electrons between nonmetals.

Naming Ionic Compounds

Naming involves the cation name followed by the anion name, adjusted for charge.

Polyatomic Ions

Ions that consist of more than one atom.

Oxyanions

Polyatomic ions containing oxygen and another element.

Covalent Compounds Naming

Use prefixes to indicate the number of atoms when naming covalent compounds.

Acids

Molecular compounds that release H+ ions in solution.

Electronegativity

A measure of an atom's ability to attract electrons.

Binary Acids

Acids that consist of only two elements, one of which is hydrogen.

Oxyacids

Acids that contain hydrogen, oxygen, and another element.

Molecular Geometry

The three-dimensional arrangement of atoms in a molecule.

Polarity of Molecules

Determined by the uneven distribution of electron density within a molecule.

Hydrogen Bonding

A weak attractive interaction between a hydrogen atom and an electronegative atom.

London Dispersion Forces

Weak attractions between all molecules due to temporary dipoles.

Ion Dipole Interactions

Attraction between an ion and the partial charge on a polar molecule.

Surface Tension

The tendency of liquid surfaces to shrink to minimize surface area.

Viscosity

A measure of a fluid's resistance to flow.

Solubility

The ability of a substance to dissolve in a solvent.

Concentration

The amount of solute per unit volume of solution.

Reaction Rate

The speed at which reactants are converted into products.

Catalysts

Substances that speed up a chemical reaction without being consumed.

Equilibrium

The state in a reversible reaction where the rates of forward and reverse reactions are equal.

Le Chatelier's Principle

If a system at equilibrium is disturbed, the system shifts to counteract the disturbance.

Electrochemical Cells

Devices that convert chemical energy into electrical energy through redox reactions.

Anode

The electrode where oxidation occurs in an electrochemical cell.

Cathode

The electrode where reduction occurs in an electrochemical cell.

Standard Reduction Potential

The tendency of a chemical species to be reduced, measured in volts.

pH Scale

A measure of the acidity or alkalinity of a solution.

Buffer Solutions

Solutions that resist changes in pH when small amounts of acid or base are added.

Redox Reactions

Chemical reactions involving the transfer of electrons between species.

Oxidation State

The charge an atom would have if all bonds were 100% ionic.

Stoichiometry

The calculation of reactants and products in chemical reactions.

Atomic Mass Unit (amu)

A unit of mass used to express atomic and molecular weights equivalent to 1.66 × 10^-27 kg.

Mass Number

The total number of protons and neutrons in an atomic nucleus.

Electron Configuration

The distribution of electrons among the orbitals of an atom.

Octet Rule

The principle that atoms tend to bond in such a way that they have eight electrons in their valence shell.

Diatomic Molecules

Molecules made up of two atoms, either of the same or different elements.

Molarity (M)

A way of expressing concentration, defined as moles of solute per liter of solution.

Leptons

A family of elementary particles that includes electrons and neutrinos.

Quarks

Elementary particles that combine to form protons and neutrons.

Chemical Reaction

A process that transforms one or more substances into different substances.

Endothermic Reaction

A reaction that absorbs energy from its surroundings.

Exothermic Reaction

A reaction that releases energy, usually in the form of heat.

Atomic Mass

The weighted average mass of an element's isotopes.

Chemical Elements

Substances that cannot be broken down into simpler substances, consisting of atoms with the same number of protons.

Molecules

Two or more atoms chemically bonded together.

Compounds

Substances formed when two or more different elements bond together.

Chemical Bonds

Forces that hold atoms together in molecules or compounds.

Ionic Bonds

Type of bond formed through the electrostatic attraction between oppositely charged ions.

Covalent Bonds

Bonds formed by the sharing of electrons between nonmetal atoms.

Mixtures

Physical combinations of two or more substances that retain their individual properties.

Homogeneous Mixtures

Uniform mixtures where the composition is consistent throughout.