Unit 2: Chemical Bonding Test Review

Vocabulary flashcards reviewing key concepts, definitions, rules, formulas, and periodic trends from Unit 2: Chemical Bonding.

Ionic bond

A bond formed between a cation and an anion (or metals and nonmetals) where electrons are transferred from the metal to the nonmetal.

Non-polar covalent bond

A chemical bond where two non-metals share valence electrons equally, having an electronegativity difference ($\Delta\text{EN}$) of less than $0.5$.

Polar covalent bond

A chemical bond where two nonmetals share valence electrons unequally, having an electronegativity difference ($\Delta\text{EN}$) of $0.5$ and above.

Metallic bond

A bond between metal atoms characterized by a "sea of electrons" where valence electrons move freely throughout the metal structure.

Electronegativity

The ability of an atom to attract shared valence electrons.

Ionization energy

The energy required to remove one valence electron from an atom.

Atomic radius (atomic size)

The distance from the nucleus to the valence electrons (the edge of the electron cloud).

Lone pair of electrons

Non-bonding valence electrons located on an atom.

Valence electrons

Electrons located in the outermost shell that are available to bond.

Cation

A positively charged ion formed when an atom loses electrons.

Anion

A negatively charged ion formed when an atom gains electrons.

VSEPR

Valence shell electron pairs repulsion theory, used to predict the shape and bond angle of a molecule based on electron repulsion.

Octet rule

The principle that atoms need $8$ valence electrons to be stable.

Lewis Dot Structure

A representation of how valence electrons are bonded and arranged in a covalent molecule.

Electrostatic attraction

The force where positive attracts negative charge, holding an ionic bond together.

Physical properties of ionic compounds

Hard, brittle, possess a lattice structure, have high melting and boiling points, and conduct electricity in solution.

Physical properties of metallic bonds

Malleable, ductile, good conductors of heat and electricity, and have high melting and boiling points.

Physical properties of covalent compounds

Have low melting and boiling points compared to ionic and metallic bonds, and are poor conductors of electricity.

Binary ionic compound naming rule

The metal keeps its name as is on the Periodic Table, and the non-metal ending is changed to -ide.

Polyatomic ionic compound naming rule

The metal keeps its name as is on the Periodic Table, and the polyatomic ion retains its specific name from the reference sheet (such as -ite or -ate).

Covalent compound naming rule

Use prefixes for both elements (except when the first element only has $1$ atom), keep the first element name from the Periodic Table, and change the second element ending to -ide.

Molecular geometry of $CO_2$

Linear shape with a $180^\circ$ bond angle because there are no lone pairs on the central carbon atom and only $2$ electron regions.

Molecular geometry of $H_2O$

Bent shape (Bent 2) with a bond angle of less than $109.5^\circ$ (or $\sim 105^\circ$) because the central atom has $2$ lone pairs that take up space.

Molecular geometry of $HCN$

Linear shape with a $180^\circ$ bond angle.

Molecular geometry of $PF_3$

Trigonal pyramidal shape with a bond angle of $107^\circ$.

Molecular geometry of $NO_2^-$

Bent shape (Bent 1) with a bond angle of $116^\circ$.

Molecular geometry of $SCN^-$

Linear shape with a $180^\circ$ bond angle.

Stable ion formation of Aluminum

Loses $3$ electrons to become a cation with a $+3$ charge ($Al^{3+}$).

Stable ion formation of Rubidium

Loses $1$ electron to become a cation with a $+1$ charge ($Rb^+$).

Stable ion formation of Bromine

Gains $1$ electron to become an anion with a $-1$ charge ($Br^-$).

Stable ion formation of Phosphorus

Gains $3$ electrons to become an anion with a $-3$ charge ($P^{3-}$).

Stable ion formation of Oxygen

Gains $2$ electrons to become an anion with a $-2$ charge ($O^{2-}$).

Stable ion formation of Strontium

Loses $2$ electrons to become a cation with a $+2$ charge ($Sr^{2+}$).

Sodium bromide

An ionic compound with the chemical formula $NaBr$.

Calcium iodide

An ionic compound with the chemical formula $CaI_2$.

Potassium oxide

An ionic compound with the chemical formula $K_2O$.

Aluminum sulfide

An ionic compound with the chemical formula $Al_2S_3$.

Hydrogen phosphate

A chemical compound with the formula $H_3PO_4$.

Ammonium nitrate

A polyatomic ionic compound with the chemical formula $NH_4NO_3$.

Strontium phosphite

A polyatomic ionic compound with the chemical formula $Sr_3(PO_3)_2$.

Dinitrogen pentoxide

A covalent compound with the chemical formula $N_2O_5$.

Carbon tetrachloride

A covalent compound with the chemical formula $CCl_4$.

$$Ca(NO_2)_2$$

An ionic compound with the chemical name Calcium nitrite.

$$KNO_3$$

An ionic compound with the chemical name Potassium nitrate.

$$Li_2O$$

An ionic compound with the chemical name Lithium oxide.

$$Li_3PO_4$$

A polyatomic ionic compound with the chemical name Lithium phosphate.

$$Li_3PO_3$$

A polyatomic ionic compound with the chemical name Lithium phosphite.

$$Li_3P$$

A binary ionic compound with the chemical name Lithium phosphide.

$$N_2O_4$$

A covalent compound with the chemical name Dinitrogen Tetraoxide.

$$SO_2$$

A covalent compound with the chemical name Sulfur dioxide.

Atomic radius group trend

Increases from top to bottom down a group because additional energy levels (shells) are added.

Atomic radius period trend

Increases from right to left across a period (decreases left to right) because more protons to the right produce more nuclear charge and more force, pulling valence electrons in closer.

Ionization energy group trend

Increases from bottom to top up a group because atoms at the top are smaller with less distance between the nucleus and valence electrons, resulting in more force that makes electrons harder to remove.

Ionization energy period trend

Increases from left to right across a period because atoms are smaller and have more protons, meaning more charge and less distance exert greater force, making valence electrons harder to remove.

Electronegativity periodic trends

Increases from bottom to top of a group and from left to right across a period because smaller atoms and higher nuclear charge create less distance and more force to attract shared valence electrons.