ELECTRONIC STRUCTURE & PERIODIC PROPERTIES OF ELEMENTS

1A

alkali metals

2A

alkaline earth metals

B

transition metals

5A

Pnictogens

6A

Chalcogens

7A

Halogens

8A

Noble gases

Metals

good conductors of heat and electricity \n malleable \n ductile \n shiny \n most are solid \n except murcury

Nonmetals

poor conductors of heat and electricity

many colors \n solids, liquids, gases \n better insulators

Wavelength

The distance between two corresponding parts of a wave

Frequency

(v) how often a new wave arrives at a point

Amplitude

(A) height of wave from midline to peak or trough

relationship of wavelength and frequency

inversely proportional as one goes up the other goes down

c= (landa)\* V

Constructive interference

The interference that occurs when two waves combine to make a wave with a larger amplitude

Destructive interference

The interference that occurs when two waves combine to make a wave with a smaller amplitude

Quantization

only those waves having an integer, n, of half-wavelengths between the end points can form

Nodes

one or more points between the two end points that are not in motion

where amplitude=0

nodes-1

radial nodes

all angles, constant radii (circles)

angular nodes

all radii, constant angles (lines)

Blackbody radiation

higher temperatures lead to greater intensity and peak shifting to shorter wavelengths

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but there is a steep drop-off in the ultraviolet region

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The peaks are in infrared at lower temps and peaks at ultraviolet/visible at increase temp

The photoelectric effect

light that meets certain energy/frequency requirements is able to eject electrons from a metal surface

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indicates that light can act like particles called photons

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the energy of photons depends on frequency

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E=hv

Line Spectra

the light emission only at specific wavelengths (because its quantized)

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Wavelength can be calculated for each of the lines in the line spectrum using the Rydberg equation

The Bohr Model

explains the line spectrum of the hydrogen atom

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In his theory of the atom, electrons are found in "orbits" around the nucleus

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Orbits that are farther out are higher in energy

Absorption

an electron moves to higher energy as light is positive

Emission

electrons moves to lower energy as light is negative

Ground state

The lowest energy state of an atom

Excited state

A state in which an atom has a higher potential energy than it has in its ground state

Rydberg Formula

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Development of quantum theory

particles and waves behave very differently on a macroscopic scale

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there is no clear separation between particles and waves

DeBroglie wavelength

λ\=h/mv

DeBroglie Mass

m\=h/λv

Heisenberg Uncertainty Principle:

we can't know exactly where an electron is, but we can predict regions where it is most likely to be found

Electron interference patterns

wave-particle duality seen with electron interference patterns

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with a few electrons, you can see the individuality of the particles

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with many electrons, you can see the wave-like patterns

Orbital

the region where an electron is likely to be found

Schrodinger equation

mathematically describes the orbitals in a hydrogen atom

Quantum numbers overview

tells us the size, shape, and orientation of the orbital each electron is in

Principal quantum \#

n, size/energy level

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corresponds to the n # in the Bohr's model

angular momentum quantum number

symbolized by l, indicates the shape of the orbital

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0-s, 1-p, 2-d, 3-f

magnetic quantum number

symbolized by m, indicates the orientation of an orbital around the nucleus

Electron spin quantum number

(ms)

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Can be +1/2 (spin up) or -1/2 (spin down)

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Related to whether electron is spinning clockwise or counterclockwise

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Need n, l, ml , and ms to describe a particular electron within an orbital

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There can be up to 2 electrons per orbital

How to find the \# of radial nodes

n-L-1

S

L- 0

ml- -1,0,1 \n orbitals- s-1 \n spheres

P

L- 1 \n ml- -1,0,1 \n orbitals- px,py,pz -3 orbitals \n dumbbells

d

L- 2 \n ml- -2,-1,0,1,2 \n orbitals- dxy, dxz, dyz, dz2, dx2,- 5 orbitals \n clovers

f

L- 3 \n ml- -3,-2,-1,0,1,2,3 \n orbitals- 7 different f orbitals

Pauli Exclusion Principle

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

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each set of 4 quantum numbers describes one electron just like each address describes how to find one person

Aufbau Principle

Fill in electron configurations and orbital diagrams starting with lowest energy orbitals

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whenever you have a choice between a lower energy orbital and a higher energy orbital, choose the lower energy one

Hund's rule

for orbitals of the same energy, maximize the number of unpaired electrons

Valence

outer shell (highest n) electrons

Core

inner (lower n) electrons

Exceptions:

Chromium to make half-full d subshell (4s1 3d5) and Copper to make full d subshell (4s1 3d10)

Electron configuration with anion

add electrons using filling rules

electron configuration with cation

take away electrons from the highest n first

Isoelectronic atoms and ion

when they have the same number of electrons and electron configuration

effective nuclear charge (Zeff)

amount of positive charge from the nucleus that an electron "feels"

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z (atomic number) - shielding (core electrons do the shielding)

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Core electrons shield better than outer electrons

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Increases across the period because core electrons remain the same but Z increases

Trend in radius

larger when moving down and left

The less protons and more electrons sticking out will increase radius \n \n The more protons than electrons the more the electrons will be pulled in by protons and radius will decrease \n \n As # of energy levels increases, electrons are farther from the nucleus and feel less attractive forces (less pull inwards by protons)

Groups

columns on the periodic table

Periods

rows on the periodic table

Ionic radius anion

\-Adding electrons makes radius bigger (more electron-electron repulsions) anion

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larger than neutral bc of electron-electron repulsions

Ionic radius cation

\-Removing electrons makes radius smaller (easier for protons to pull electrons inwards) cation

\-Removing all valence electrons makes it MUCH smaller (going from larger n to smaller n) \n - smaller than neutral

Ionic radius trend

increases as you go down and left

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\-For isoelectronic series (different # of protons but same # of electrons), radius decreases with increasing Z

Ionization energy

The amount of energy required to remove an electron from an atom in the gas phase

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the result is a cation with a positive charge

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upper right= higher number = harder to remove

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lower left=lower number=easier to remove

Exceptions of ionization energy trend (3A)

Ionization energy slightly decreases from 2A to 3A

Exceptions of ionization energy trend (6A)

ionization slightly decreases from 5A to 6A because it is easier to remove a paired electron from the 6A

Why does it take more energy to remove core electrons than valence electrons?

because they feel more attraction to the nucleus and there is less shielding

Which electron is the hardest to remove when having 4 valence electrons?

the 4th IE4

Electron Affinity

the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion (anion)

Electron affinity trend

increases as it goes to upper right but the exception is the noble gases because they do not want electrons

Ionic compounds

3D array of ions, made of cations and anions held together by electrostatically

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they can make single elements or multiple elements

Molecular compounds

covalent bonded compounds

Monoatomic

single element ion compound

Polyatomic

multiple atoms of ionic compounds

What are metals most likely to become and why?

cations because they tend to lose electrons and they have low ionization energy so that they can have the same number of electrons as a noble gas

What are non metals most likely to become and why?

anions because they have high electron affinities

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This is so they have the same number of electrons as a nobel gas

Cations charge

charge \= group \#

anions charge

charge\=group \#-8

Perchlorate

ClO4-

Chlorate

ClO3-

Chlorite

ClO2-

Hypochlorite

ClO-

nitrate

NO3-

nitrite

NO2-

sulfate

SO4 2-

sulfite

SO3 2-

phosphate

PO4 3-

phosphite

PO3 3-

perchloric acid

HClO4

Chloric acid

HClO3

chlorous acid

HClO2

hypochlorous acid

HClO

nitric acid

HNO3

nitrous acid

HNO2

sulfuric acid

H2SO4

sulfurous acid

H2SO3

phosphoric acid

H3PO4

Phosphorous acid

H3PO3

Hydronium

H3O+

water

H2O

hydroxide

OH-

Sulfate

SO4 2-