CHEM 20 Chapter 1: Structure & Bonding, Acids & Bases

Vocabulary flashcards created from Chapter 1 Study Guide covering atomic structure, hybridization, geometry, electronegativity, and acid-base concepts.

Proton

A subatomic particle located in the nucleus with a charge of $+1$.

Neutron

A subatomic particle located in the nucleus with a charge of $0$.

Electron

A subatomic particle located outside the nucleus with a charge of $-1$.

Atomic number

The total number of protons in an atom. Z is the symbol used to represent it.

Mass number

The total number of protons plus neutrons in an atom. A is the symbol.

Valence electrons

Electrons in the outermost shell that control bonding and chemical reactivity.

s orbital

A spherical atomic orbital that holds a maximum of $2$ electrons.

p orbital

A dumbbell-shaped atomic orbital that holds a maximum of $2$ electrons.

Hydrogen

1 valance e

1 bond

0 lone pairs

Carbon

4 valance e

4 bonds

0 lone pairs

Nitrogen

5 valance e

3 bonds

1 lone pairs

Oxygen

6 valance e

2 bonds

2 lone pairs

Lewis structure steps

1. count total valence electrons

2. connect atoms with single bonds

3. complete outer shell

4. place remaining electrons as lone pairs

Duet Rule

Hydrogen follows the duet rule (2 electrons ) not the octet rule

Sigma (sigma) bond

A covalent bond formed by the head-on overlap of orbitals.

Pi (pi) bond

A covalent bond formed by the sideways overlap of unhybridized p orbitals.

single pond

= 1 sigma bond

double bond

= 1 signma bond + 1 pi bond

triple bond

= 1 sigma bond and 2 pi bonds

$sp^3$ hybridization

Hybridization involving $4$ electron groups, tetrahedral geometry, bond angles of $109.5^\circ$, and single bonds.

$sp^2$ hybridization

Hybridization involving $3$ electron groups, trigonal planar geometry, bond angles of $120^\circ$, and double bonds.

$sp$ hybridization

Hybridization involving $2$ electron groups, linear geometry, bond angles of $180^\circ$, and triple bonds or two double bonds

Electronegativity

An atom's ability to attract shared electrons in a bond.

Electronegativity patter

Top, right of periodic table is most electronegative

-more electronegative atom carries partial negative charge and less electronegative atom carries partial positive charge

pka and electronegativity

lower pka means higher electronegativity

Ionic bonds

electron transfer permanently to another atom

change in EN > 2

Covalent bonds

electrons shared equally

change in EN < 0.5

Polar bonds

Electrons are not shared equally

change in EN= 0.5-2.0

Brønsted–Lowry acid

$H^+$ donor.

Brønsted–Lowry base

$H^+$ acceptor.

Conjugate acid/base pairs relationship

stonger acid <> weak conj base

weaker acid <> strong conj base

predicting equilibrium direction

1. use pka to guide you

2. H+ goes from stronger acid to stronger base

3. Acid with smaller pka will give up H+ to conj base of acid with larger pka.

Lewis acid

An electron-pair acceptor.

Lewis base

An electron-pair donor.

Lone pair in lewis acid-base reaction

can act as the electron source

Curved arrow

A notation that shows the movement of an electron pair, always starting at the electrons and pointing to where they move.