Unit 1: Matter, Chemical Trends, and Chemical Bonding

Comprehensive vocabulary flashcards covering atomic theory, periodic trends, chemical nomenclature, bonding models, and the physical properties of ionic and molecular compounds.

Atomos

A Greek term meaning "uncuttable," used by ancient Mediterranean philosophers to describe the indivisible building block reached if an object were repeatedly cut in half.

Billiard Ball Model

John Dalton's 1803 model of the atom, which conceptualized atoms as solid, indivisible spheres that are unique for each element.

Plum Pudding Model

J.J. Thomson's 1897 atomic model consisting of a sphere of evenly distributed positive charge with negatively charged electrons embedded within it.

Planetary Model

Ernest Rutherford's 1911 model of the atom in which planet-like electrons orbit a tiny, massive, positively charged central nucleus.

Bohr Model

An atomic model proposed by Niels Bohr in 1913 that refined Rutherford's model by establishing that electrons can occupy only discrete, allowed energy levels or shells.

Electron Cloud Model

The quantum mechanical atomic model based on Erwin Schrödinger's 1926 wave equation, describing regions of space around the nucleus where electrons are most likely to be located.

Valence Electrons

Electrons located in the outermost shell of an atom that are exclusively involved in the formation of chemical bonds.

Lewis Diagram

A model of an atom that features the chemical symbol of the element surrounded by dots representing its valence electrons.

Electron Pairs

Two valence electrons in the same shell that interact in a unique way allowing them to be situated close together, making them less likely to participate in bond formation.

Unpaired Electrons

Electrons in an unfilled outer shell that do not form a pair and are therefore available and likely to participate in chemical bonding.

Strong Nuclear Force

the force holding protons and neutrons together, approximetly 40 times greater than the repultion betweeen protons

Isotopes

Atoms of the same chemical element that contain the same number of protons but different numbers of neutrons.

Atomic Number ($Z$)

The number of protons in the nucleus of an atom, which specifies the identity and chemical nature of an element.

Mass Number ($A$)

The total sum of the number of protons and the number of neutrons contained within the nucleus of an atom.

Radioisotopes

Isotopes with unstable nuclei that spontaneously decay into different, often stable, nuclei by emitting particles or energy.

Alpha
Beta
Gamma

Atomic Mass Unit ($\text{u}$)

A standard unit of mass defined as exactly one-twelfth of the mass of a carbon-12 atom, equal to $1.66 \times 10^{-24}\,\text{g}$.

Isotopic Abundance

The percentage of a particular naturally occurring isotope of an element relative to the total quantity of all isotopes of that element found on Earth.

Electron Configuration

the arrangement of electrons in the energy levels and orbitals of an atom

Period

A horizontal row in the periodic table corresponding to the number of occupied electron shells in the atoms of those elements.

Group

A vertical column in the periodic table containing elements with identical valence shell electron configurations and similar chemical properties.

Main-Group Elements

The elements in Groups 1, 2, and 12 through 18 whose valence electron counts are generally predictable from their column positions.

Transition Metals

Elements located in Groups 3 through 11 that are hard metals with high melting points and complex electron configurations capable of holding up to 18 electrons in their outer shell. (dont follow the octet rule)

Inner Transition Elements

The two rows of metallic elements at the bottom of the periodic table, consisting of lanthanoids and actinoids, that fit between Groups 3 and 4.

Alkali Metals

The metallic elements belonging to Group 1 of the periodic table (excluding hydrogen), which are highly reactive and react vigorously or explosively with water.

Alkaline Earth Metals

The metallic elements belonging to Group 2 of the periodic table that are reactive, but less reactive than the alkali metals.

Halogens

The reactive non-metal elements located in Group 17 of the periodic table, possessing seven valence electrons.

Noble Gases

The extremely unreactive elements in Group 18 of the periodic table that naturally possess completely filled valence electron shells.

Metalloids

Elements with physical and chemical properties intermediate between metals and non-metals; they are typically shiny and brittle semiconductors.

Lanthanoids

The inner transition elements of Period 6 ranging from lanthanum through lutetium, historically called rare earth elements.

Actinoids

The inner transition elements of Period 7 ranging from actinium through lawrencium, none of which possess stable isotopes and all of which are radioactive.

Atomic Radius

The distance measured from the centre of an atom to the boundary within which the electrons spend $90\%$ of their time.

Effective Nuclear Charge

the positive pull (attraction from the protons) an electron feels after shielding (the inner energy levels block out the force off attraction and electrons and electrons repel eachother)

Ionization Energy

The amount of energy required to remove the outermost electron from an atom or ion in the gaseous state.

Magent and paperclip example — my hand whould work as the ionization energy

Electron Affinity

The quantity of energy absorbed or released when an electron is added to a neutral atom in the gaseous state.

both repultion and attraction effect the change,

an electron joins the system, (an atom) —> stronger attraction and repulion —> change in energy.

• overall increase in energy — energy is absord

• overall decrease — energy is released

Electronegativity

an atom's ability to attract shared electrons to itself in a chemical bond.

Octet Rule

A rule of thumb stating that when bonds form between atoms, the atoms tend to gain, lose, or share electrons until they achieve a filled valence shell containing eight electrons.

Ionic Bond

The attractive electrostatic force that holds oppositely charged ions together after the transfer of one or more valence electrons.

Ionic Compound

A chemical compound composed of positive ions and negative ions held together by ionic bonds in ratios that produce a zero net charge.

Covalent Bond

The attractive force between two atoms resulting from the mutual electrostatic attraction of their positively charged nuclei for one or more shared pairs of electrons.

Molecular Compound

A chemical compound composed exclusively of non-metal atoms joined together by covalent bonds.

Single Bond

A covalent bond formed by the sharing of one pair of electrons ($2$ electrons) between two atoms.

Double Bond

A covalent bond formed by the sharing of two pairs of electrons ($4$ electrons) between two atoms.

Triple Bond

A covalent bond formed by the sharing of three pairs of electrons ($6$ electrons) between two atoms.

Bonding Pair

A pair of electrons that is shared between two atoms, directly constituting a covalent bond.

Lone Pair

A pair of valence electrons on an atom that is not involved in covalent bonding with another atom.

Lewis Structure

A Lewis diagram modified to represent a complete molecular compound or polyatomic ion, showing bonded atoms, bonding pairs, and lone pairs.

Polyatomic Ion

A molecular species consisting of two or more covalently bonded atoms that carries an overall positive or negative charge due to an excess or deficit of electrons.

Polar Covalent Bond

A covalent bond with an uneven distribution of shared electron density, producing a partial positive charge ($\delta+$ ) at one end and a partial negative charge ($\delta-$) at the other.

0.0 - purly covalent,

0.0 - 0.4 — non polar, (london disperion - the attraction between induced dipoles and temperarary dipoles)

0.4 to 1.7 — polar

1.7 and above — ionic

Electronegativity Difference ($\Delta\text{EN}$)

The difference between the electronegativity values of two bonded atoms, used to assess the degree of ionic or covalent character in a bond.

Structural Formula

A chemical diagram in which chemical symbols are connected by lines representing covalent bonds, omitting non-bonding lone pairs.

Alkali

A water-soluble base, historically named from the Arabic word al-qali meaning "the ashes."

Oxoacid and binary acids

oxyacid - An acid composed of hydrogen, oxygen, and atoms of at least one other element, usually a non-metal. (Sulfuric asid- H₂SO₄)

• if polyatomic ion ends on ate - add IC ACID and if it ends on ite add OUS ACID

binary - no oxygen (hydrochloric acid - HCl)

• hydro- + nonmetal root + -ic acid

Melting Point

The temperature at which a pure compound changes state from a solid to a liquid at standard atmospheric pressure ($101.325\,\text{kPa}$).

Boiling Point

The temperature at which a pure compound transitions from a liquid to a gas at standard atmospheric pressure ($101.325\,\text{kPa}$).

Dipole

A polar molecule exhibiting an asymmetrical charge distribution, resulting in a slightly positively charged pole and a slightly negatively charged pole.

Dipole-Dipole Force

The attractive electrostatic intermolecular force occurring between the positively charged pole of one polar molecule and the negatively charged pole of an adjacent polar molecule.

Intermolecular Forces

Attractive forces that act between individual molecules, including dipole-dipole attractions and weak attractions between non-polar molecules.

Electrical Conductivity

The ability of an object or substance to allow an electric current to pass through it due to the free, independent movement of charged particles such as electrons or ions.

Nonylphenol Ethoxylates (NPEs)

Surfactants — lets water interact with oily substances and dirt,

Nonylphenol — interacts with oils

• a carbon-based chemical group with a ring-shaped structure and a nine-carbon chain.

Ethoxylates — interacts with water

• repeating units added to the nonylphenol to help the chemical interact with