Atomic Structure Notes

Atomic Structure Notes

Electromagnetic Radiation (radiant energy)

§ Energy through space

§ forms of radiant energy are gamma, x rays, UV light, visible light, infared light, microwaves, radio waves

§ Visible light- light we see with our eyes (wavelength of 400-700 nm)

§ each has their own range of wavelength and frequency

§ all electromagnetic radiation moves through a vacuum at the speed of light (c) which equals 3.00 x 108 m/s.

§ electromagnetic spectrum displays all forms of radiant energy in terms of increasing wavelength

gamma Xray UV visible IR microwaves Radio

à à à à à à à wavelength increases

ß ß ß ß ß ß ß frequency decreases

Parts of a Wave

§ wavelength (λ- lambda)– the distance between to successive peaks or troughs; usually measured in nanometers (1nm = 1 x 10-9 m) or micrometers (1 um = 1x 10-6 m ) 1 A = 1x 10-10 m

§ frequency (ν- nu) – the number of cycles that pass a given point each second; measured in hertz ( 1 Hz = 1 cycle per second or 1/s or s-1)

§ wavelength and frequency are inversely proportional to one another

c= λ ν

Examples 1-2

Quanta and Photons

§ Matter and Energy were seen as different from each other

Matter à made of particles Energy à made of waves

§ Max Planck (black body radiation)

1. when solids are heated, they release radiation with various intensity and color based on their temperature

2. assumed energy can only be absorbed or released in “chunks”, “packets”, or “quanta”

with the size of hν (whole number multiples)

§ E = h ν

§ h (planck’s constant) = 6.63 x 10-34 J-s

§ Energy is proportional to frequency

If that is true, why do energy changes seem continuous?

On the macroscopic level, hv is such a small amount it goes unnoticed

On the atomic level, the impact is much more significant

§ Albert Einstein (photoelectric effect)

1. Light at a certain minimum frequency shown on a metal surface emitted electrons

2. used Planck’s quantum theory to explain it; assumed the radiant energy striking the metal surface was a stream of tiny energy packets which he termed a “photon”. A photon behaves like a particle.

3. so the E in E = hν corresponds to the energy of a photon

E = h ν so E = hc/ λ if we assume the E in E=mc2 is the same energy just mentioned we can get an equation for the apparent mass of a photon

m = h/ λ c

(helps us explain high frequency(short wavelength) of X rays giving photons the ability to cause damage )

§ So which is it? Is energy a wave or a particle?

§ It’s both! The concept is called wave-particle duality. Light on the macroscopic level exists as waves but on the microscopic level, it consists of a collection of photons

§ What about the other way?

§ Can electrons which are particles behave like waves? See De Broglie section

Line Spectra and the Bohr Model

§ Radiation of a single wavelength (laser) is monochromatic.

§ Most radiation sources contains many different wavelengths producing a spectrum when separated by a prism

§ Continuous spectrum- white light- all of the colors seen (roy g biv)

§ Line Spectrum- specific gases emit different colors at specific wavelengths-like a fingerprint

§ Line spectrum of hydrogen ( 4lines)

§ Used Rydberg equation to calculate the wavelengths of the 4 spectral lines of hydrogen 1/ λ = (RH ) ( 1/(n1) 2 – 1/(n2) 2 ) Rydberg Constant = 1.10 x 107 m-1

Bohr Model tried to explain line spectrum of hydrogen

§ Assumed electrons moved in circular orbits like the planets around the sun

§ Used Planck’s idea and assumed electrons stay within the orbits because they have a fixed, definite energy for specified energy levels(orbits)

§ Also assumed that energy is only released or absorbed by an electron as it changes from one energy level to another; an electrons cannot reside between energy levels

1. To move to a higher energy level, electrons must absorb energy

2. To fall back down to a lower energy level, electrons release energy

Energy states of the hydrogen atom

§ Energy levels close to the nucleus are of less energy but are more stable; used n to symbolized the energy level number n =1, 1st energy level

§ Ground state- lowest energy state of an atom

§ Excited state- when a electron is in an higher energy level than its most stable state

§ Calculate the energy of specific energy level E= (-2.18 x 10-18 J)(Z2/n2)

n = energy level

Z= charge of the nucleus

§ Modifying the equation to show when an electron moves from one energy level to another, ΔE = Ef – Ei

§ So ΔE= (-2.18 x 10-18 J)(1/nf 2 - 1/ni 2)

§ His equation matched the Rydberg equation very well

§ A negative number means energy is being released

Limitations of Bohr model

§ It could not explain the spectra of other atoms

§ Electrons do not circle the nucleus

§ The ideas we keep are:

1. Electrons exist only in certain energy levels which will be described by quantum numbers

2. Energy is released or absorbed as it jumps from one level to another

3. Ground state and excited state terms

DeBroglie’s Discovery

§ Assumed if light waves had particulate behavior, then particles like electron can have wave-like properties

§ The equation: λ = h/ mv Don’t mistake the v of velocity for ν (nu)

Predicts that matter with normal mass creates small waves but matter with extremely small mass , like an electron traveling at high speed emit decent size wavelengths.

Bohr Model OUT, Quantum Mechanical Model IN

§ New approach – represented by a complex mathematical equation. The solutions described the energies and spatial distribution of electrons

§ In this model:

a. we don’t know the exact location, we can only give a probability of finding an electron within an atom within an orbital

b. orbital- 3-dimensional space where there is a high probability of finding an electron

c. two ways of showing an atomic orbital:

1. electron cloud- the high density reflects the probability of finding the electron

2. a surface that contains 90% of electron probability

d. Heisenburg Uncertainty Principle tells us there is a limit to how well we can know the position and momentum of an object

§ Schrodinger’s equation’s solutions (many wave functions) are described by quantum numbers

Quantum Numbers

§ Quantum numbers describe the properties of electrons within atomic orbitals

1. Principal quantum number (n)- indicates the main energy level(shell) occupied by an electron

a. values of n = 1-7

b. as n increases, the electron’s energy increases and so does its distance from the nucleus.

c. Maximum number of electrons within an energy level 2 n2

2. Angular momentum quantum number (l) indicated the shape of the

orbital(subshell)

a. values ranges from 0 to (n-1)

b. l =0 ( s) spherical ; l = 1 (p) dumbbell ; l = 2 (d) four lobed; l = 3 (f) too complex; l =4 (g)….

c. in the nth energy level there are n sublevels

3. magnetic quantum number (ml) – indicates the orientation of an orbital around

the nucleus (how many orbitals of a sublevel)

a. values range from – l to + l including zero

b. s= 1 orientation; p= 3 orientations; d= 5 orientations; f=7 orientations

c. Total number of orbitals within energy level is n2

d. All orbitals within the same energy level are degenerate – same energy

4. Electron Spin quantum number (ms)- show the electron moving in a

counterclockwise or clockwise direction

a. can only have 2 values (+ ½ or – ½ )

SO ….

1st energy level s orbital only 1 shape

2nd energy level p orbitals 3 shapes

3rd energy level d orbitals 5 shapes

4th energy level f orbitals 7 shapes

Three Rules to Electron Placement

1. Aufbau Principle: electrons fill orbitals from the lowest to highest energy level

§ Use the periodic table to guide you: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 4d…….increasing energy à

§ Complete electron configuration for :

§ Noble Gas (shorthand) configuration; place the noble gas that comes before your element in brackets and start your configuration from the noble gas

2. Hund’s rule : lowest energy configuration has the maximum number of unpaired electrons

§ Fill each orbital of equal energy with one electron before pairing

§ Orbital notation- use lines, boxes or circles to represent orbitals and arrows to represent the electrons within the orbitals

3. Pauli exclusion Principle- no atom can have the same 4 quantum numbers so

an orbital can only hold 2 electrons and they must be of opposite spins (↑↓)

§ Valence electrons- electrons located in the highest energy level(outer electrons within the s and p orbitals)

§ Core electron- inner shell electrons

Special Configurations(anomalies)

§ Not thoroughly explained but within the transitions and inner transitions, there are variances

§ The Aufbau principle can be used to write the electron configurations for all atoms up to V on the periodic chart and then variances occur throughout the rest of the table

§ BECAUSE

1. Completely filled d sublevel is more stable than a half filled or partially filled sublevel

2. Half filled d sublevel is more stable than a partially filled sublevel

§ Know the anomalies for Cu, Cr, Mo, Ag, and Au

Magnetism

§ Paramagnetic – electrons are unpaired so they are attracted to a magnetic field

§ Diamagnetic – electrons are paired so they are repelled by a magnetic field

Ions

§ Charged particles

§ Cations(+ ions)- metals, lose electrons to acquire 8 electrons in their outer shell

§ Anions(- ions) – nonmetals, gain electrons to acquire 8 electrons in their outer shell

§ In the S and P block (group A representative metals)

1. metals – group 1A – 3A, group number is the charge number

2. nonmetals- groups 4A- &A, group number –8 is the charge number

§ D block ( group B transition metals)

1. when forming positive ions, they lose the outer s electrons first before any d electrons

§ Isoelectronic- atoms that have the same electron configuration

The Periodic Table

Periodicity- recurring pattern of similar properties

1865 Mendeleev and (Meyer) – arranged elements known at time by increasing atomic mass noting similar chemical & physical properties of elements recurring periodically

· Mendeleev given most credit

· Predicted elements to fall in blank spaces put into the table

1913 Moseley- worked with Rutherford to develop atomic number concept; arranged the table in increasing atomic number and noted similar physical and chemical properties recurring periodically

· Filled holes in table from Mendeleev

Periodic Law - when elements are arranged in order of increasing atomic number, elements with similar properties fall into the same column

Things to remember

§ Elements in the same column has the same valence electrons and thus the same electron configuration( n will change)

§ Valence electrons determine chemical properties and some physical

Organization of the table

§ Period - horizontal row

§ Group or family – vertical column

3 classes of elements

· Metals- three quarters of table, left side, exhibit shiny luster, conduct heat and electricity, are malleable(pounded into thin sheets) and ductile(drawn into fine wire), all are solids at room temperature except mercury, have low ionization energies, high tensile strength, mostly silver in color, some colored; metals with a nonmetal form ionic solids

· Nonmetals- far right side of table, not lustrous, poor conductors of heat and electricity, generally lower melting points than metals, mixture of brittle and soft solids, gases and 1 liquid at room temperature, nonmetals with pother nonmetals form covalent(molecular) in gas, liquid or low-melting solids

· Metalloids(semimetals)-fall on the border of the staircase line with exception of aluminum and polonium, properties of both metals and nonmetals, semiconductors, all solids at room temp

Example- silicon- looks like a metals but is brittle like a nonmetal

Representative elements (main-group) - group A elements

§ S and p sublevels

· Represent similar physical and chemical properties that change in a regular way

· Columns are referred to by a roman numeral and letter-antiquity , or by numbers 1-18 or by family name

· Antiquity- group number of letter A elements tell how many valence electrons present- all elements in a column have same valence number; also denotes ion charge(groups 1A-3A; group number = charge number)(group 4A – 8A: 8- group number = charge number)

· Valence electrons primarily determine the properties of an atom

Alkali Metals- group 1A

· Soft, easily cut with knife

· Silvery, metallic luster

· Most chemically reactive metals, reactivity increases down a group

· Not found free in state

· React with nonmetals to form salts(ionic compounds)

· Reaction with water is highly exothermic and extremely volatile producing hydrogen gas and an hydroxide base

· Usually stored under kerosene or oil

· High thermal and electrical conductivity

· Forms +1 ions when the one valence electron is lost

· Low density, increases down a group and melting point decreases down a group

Alkaline Earth metals – group 2A

· Harder, and denser than 1A

· Silvery, metallic luster

· Forms metal hydroxides-thus basic or alkaline

· Extracted from mineral ores- rocks mined for minerals- thus earth

· High thermal and electrical conductivity

· Not as reactive as 1A but still chemically reactive metals, reactivity increases down a group

· Not found free in state

· React with nonmetals to form salts (ionic compounds)

· Forms +2 ions when 2 valence electrons are lost

· Higher melting points than 1A