Chapter 3: Electronic Structure + Periodic Properties of Elements

wavelength (λ)

distance between 2 crests (meters or nanometers)

frequency (𝜈)

number of cycles per second (1/s, S^-1, Hz)

speed of light in a vacuum

c = 3.00 × 10⁸ m/s

Planck

-equation: E=hv

-constant: 6.262 × 10^-34J.s

energy

capacity to do work

work

force applied through distance

force

any push or pull on an object

joule

SI unit of energy, 1J = kg (m/s)²

photoelectric effect

light shining on the surface of metal can cause electrons to be ejected (only ejected if the photons have sufficient energy)

photons

tiny packets or particles of light

Bohr’s Model

nucleus in the middle, electrons surrounding in circular orbits with quantized energy levels

Equation for energy of an electron in orbit

E = -k/n²

What is the value of the constant “k”

2.18 × 10^-18J

What happens to the energy of an electron as it gets further from the nucleus?

it decreases

de Broglie Equation

wavelength = h/ mv

constructive interference

two wavelengths in the same period combine to increase amplitude

destructive interference

waves in opposite periods combine to create a flat line with no amplitude

electron defraction

means that electrons do have wavelike tendencies and the orbits are not perfectly circular

Heisenberg uncertainty principle

it is impossible to know the exact momentum (mv) of an electron and its exact location at the same time

Probability density

represents the probability that an electron will be in that exact predicted location

Electron density

regions where there is a high probability of finding an electron (decreases further from nucleus)

Nodes

places where there is zero chance you will find an electron (as n increases, these increase)

Orbitals

wave functions of an electron in an atom, a mathematical function of their wavelike behavior

Principal #

n: the level of orbital, as n increases, the orbital becomes larger and energy increases

Azimuthal #

l: have values from 0 to 1 for each n value, this number gives the orbital shape

Magnetic quantum #

m_l: have values between -l and l, this number describes the orientation of an orbital in space

Shell

orbitals with the same value of n

Subshell

a set of orbitals with the same n and l values

S orbital

spherically symmetrical, as n gets bigger the orbital size increases

P orbital

not symmetrical, has a node at the nucleus (on the axis), dumbbell shaped

D orbitals

four leaf clover, lie in planes, d₂² has a donut shape around a p orbital, 2 nodes (one on each axis)

Pauli exclusion principle

no two electrons can have the same set of four quantum numbers

Shielding

electrons in the atom block each other from the nucleus

Effective nuclear charge

Z_eff : the net positive charge attracting the electron (as l increases, this value decreases)

Degenerate

orbitals in a sub-shells that have the same energy

Orbital diagram

a box represents the orbital and a half arrow represents each electron (paired one up and one down = ground state)

Stable electron

when they are in the lowest possible energy state in the lowest orbital

Paramagnetism

parallel or unpaired spins; strongly attracted by a magnet

Diamagnetism

antiparallel or paired spins; slightly repelled by a magnet

Hund’s rule

for degenerate orbitals, the lowest energy is attained when the # of electrons with the same spin is maximized (fill each level singly first and then double)

Octet

8 electrons in the outermost shell (ns²np⁶)

Valence electrons

electrons in the outermost shell (past the noble gas core)

Core electrons

electrons in the inner shells (takes a lot of energy to involve these electrons)

What two elements are the exception to Hund’s rule?

Chromim (Cr) and Copper (Cu): the 4s level is not paired

Periodic Law

properties of elements are periodic functions of their atomic numbers

Lewis symbol

use atomic symbol and a dot for each valence electron

Transition metal ions

lose valence electrons from the s shell first and then as many from the d shell as required to reach the charge of the ion (first in first out)

Trend of Zeff

increases left to right (because nuclear charge increases but less shielding occurs)

Trend in atomic radius

increases right to left, increases down the group

Cations

smaller than parent atom (less electron-electron repulsion)

Anions

larger than parent atoms (more electron-electron repulsion)

Trend in ionization energy

increases left to right, decreases down a group (it is easier to take away from larger more negative atoms)

Trend in electron affinities

increases left to right, increases up the groups

Trend in metallic character

increases right to left

Trend in electronegativity

increases left to right, increases up groups

Ionic compounds

generally combo of metals and non-metals (only empirical formulas can be written for most ionic compounds)

Molecular compounds

generally non-metals only