Electrons in Atoms

• Proposed that matter was composed of tiny indivisible particles

• Not based on experimental data

• Greek: atomos

• Mixture of science and mysticism.

• Lab procedures were developed, but alchemists did not perform controlled experiments like true scientists.

John Dalton (1807)

British Schoolteacherbased his theory on

others’ experimental data

Billiard Ball Modelatom is a

uniform, solid sphere

Henri Becquerel (1896)

Discovered radioactivityspontaneous emission

of radiation from the nucleus

Three types:alpha () - positivebeta () - negativegamma () - neutral

J. J. Thomson (1903)

Cathode Ray Tube Experimentsbeam of negative

particles

Discovered Electronsnegative particles

within the atom

Plum-pudding Model

J. J. Thomson (1903)

Plum-pudding Model

positive sphere (pudding) with negative electrons (plums) dispersed throughout

Ernest Rutherford (1911)

Gold Foil Experiment

Discovered the nucleusdense, positive charge

in the center of the atom

Nuclear Model

Rutherford’s Gold Foil Experiment

(a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct. (b) Actual results.

Ernest Rutherford (1911)

Nuclear Modeldense, positive nucleus

surrounded by negative electrons

Niels Bohr (1913)

Bright-Line Spectrumtried to explain

presence of specific colors in hydrogen’s spectrum

Energy Levelselectrons can only exist

in specific energy states

Planetary Model

Niels Bohr (1913)

Planetary Model

electrons move in circular orbits within specific energy levels

Bright-line spectrum

Erwin Schrödinger (1926)

Quantum mechanics electrons can only

exist in specified energy states

Electron cloud model orbital: region around

the nucleus where e- are likely to be found

Erwin Schrödinger (1926)

Electron Cloud Model (orbital) dots represent probability of finding an e-

not actual electrons

James Chadwick (1932)

Discovered neutronsneutral particles in the

nucleus of an atom

Joliot-Curie Experimentsbased his theory on

their experimental evidence

James Chadwick (1932)

Neutron Model revision of Rutherford’s Nuclear Model

Electromagnetic RadiationElectromagnetic radiation –

radiowaves, X-rays, microwaves, infrared waves, visible light, ultraviolet waves and gamma rays.

All electromagnetic radiation travel at the speed of light (c = 3.0 x 108 m/s) in a vacuum.

Physics and the Quantum Mechanical ModelAmplitude – wave’s height from the

origin to the crest.

Wavelength ()– distance between the crests.

Frequency ()– number of wave cycles to pass a given point per unit of time.

Physics and the Quantum Mechanical ModelFrequency and wavelength are inversely

proportional. As frequency increases, wavelength decreases, and vice versa, but their product will always equal the speed of light.

c =

SI units for frequency are cycles per second is a hertz (Hz), or 1/seconds (1/s or s-1).

Relationship Between Wavelength and Frequency

Physics and the Quantum Mechanical ModelWhat is the frequency of light that has a wavelength

of 550 nm? (1m = 109 nm or 1 nm = 10-9 m)?

What is the wavelength of light, in cm, that has a frequency of 9.60 x 1014 Hz (1/s)?

What is the frequency of light (Hz) that has a wavelength of 740 nm (1m = 109 nm or 1 nm = 10-9 m)?

Physics and the Quantum Mechanical ModelSunlight splits into a spectrum of colors

when it passes through a prism.

Colors of the spectrum include red, orange, yellow, green, blue, indigo and violet.

Red light has the longest wavelength and the lowest frequency, while violet light has the shortest wavelength and the highest frequency.

Dispersion of White Light By a Prism

Physics and the Quantum Mechanical ModelEvery element emits light after it absorbs energy.

The light that is emitted (atomic emission spectra) is different for every element, and differs from white light because it is not continuous.

Max Planck said that color changes can be explained if you assume that the energy of a substance changes in small increments.

Emission (line) Spectra of Some Elements

Emission (line) Spectra of Some Elements (cont’d)

Emmision (line) Spectra of Some Elements (cont’d)

Physics and the Quantum Mechanical ModelPlanck showed that the amount of radiant

energy (E) absorbed or emitted by a substance is proportional to the frequency of the radiation.

E = hh is Planck’s constant (6.626 x 10-34 J s)Any attempt to increase or decrease the

energy of a system by a fraction of h times will fail because energy is only emitted or absorbed in quanta, or bunches of energy.

Planck’s Constant ExamplesWhat is the energy of a photon with a

frequency of 2.94 x 1015 cycles per second (s-1 or Hz)?

What is the energy of a light particle with a wavelength of 675 nm?

Homework Problem ExamplesWhat is the wavelength, in nm, of light with a frequency

of 9.5 x 109 s-1? ( 1 m = 109 nm)

How much energy is contained in a photon with a wavelength of 5.17 x 10-4 m?

Planck’s RevelationShowed that light energy could be thought of

as particles for certain applications

Stated that light came in particles called quanta or photons

Particles of light have fixed amounts of energy

The energy of the photon is directly proportional to the frequency of lightHigher frequency = More energy in photons

Physics and the Quantum Mechanical ModelPhotons – light energy. The energy of

photons is quantized according to the equation E = h.

Light was therefore thought to have a dual wave-particle behavior to explain all of its characteristics.

Bohr’s ModelEnergy of an electron is related to the

distance electron is from the nucleusEnergy of the atom is quantizedatom can only have certain specific

energy states called quantum levels or energy levels

when atom gains energy, electron “moves” to a higher quantum level

when atom loses energy, electron “moves” to a lower energy level

lines in spectrum correspond to the difference in energy between levels

Bohr’s ModelAtoms have a minimum energy called the ground stateThe ground state of hydrogen corresponds to having its

one electron in an energy level that is closest to the nucleus

Energy levels higher than the ground state are called excited statesthe farther the energy level is from the nucleus, the higher its

energyTo put an electron in an excited state requires the

addition of energy to the atom; bringing the electron back to the ground state releases energy in the form of light

Bohr’s ModelDistances between energy levels decreases as

the energy increaseslight given off in a transition from the

second energy level to the first has a higher energy than light given off in a transition from the third to the second, etc.

1st energy level can hold 2 electrons (e-1), the 2nd 8e-1, the 3rd 18e-1, etc.farther from nucleus = more space = less

repulsion

Models of the AtomEnergy level – region around the

nucleus where the electron is likely to be found. Think of steps on a ladder.

Essentially, you must be in one energy level or another, you can’t be between energy levels, just like you can’t stand in mid-air between the steps of a ladder.

Models of the AtomEnergy levels are not equally spaced.

The further away an electron is from the nucleus, the easier it becomes to pull that electron off of that particular atom.

Erwin Schrodinger – in 1926, he came up with a new way of describing the energy and location of an electron, called the quantum mechanical model, which is a mathematical method.

Models of the AtomThe quantum mechanical model does

not say that electrons take exact paths around the nucleus, but that it estimates the probability (likelihood) of finding an electron in a certain position.

If the electron cloud is very dense, it is more likely that you will find the electron there, then if the electron cloud is less dense.

OrbitalsOrbital – area where an electron is likely

to be found.usually use 90% probability to set the

limitthree-dimensional

Orbitals are defined by three integer terms called the quantum numbers.

Each electron also has a fourth quantum number to represent the direction of spin

Models of the AtomPrincipal quantum number (n) –

designates the energy level of the electrons. n will always be an integer.

The distance from the nucleus increases as n increases.

Within each energy level, electrons occupy energy sublevels.

The number of energy levels (n) is always the same as the number of sublevels.

Models of the AtomSublevel – part of an energy level.

1st energy level has 1 sublevel (“s” sublevel)

2nd energy level has 2 sublevels (“s” and “p” sublevels)

3rd energy level has 3 sublevels (“s”, “p”, and “d” sublevels)

4th energy level has 4 sublevels (“s”, “p”, “d” and “f” sublevels)

Models of the AtomAtomic orbitals – areas where electrons

are likely to be found.s orbital – spherical in shape, only 1 s

orbital per sublevel.p orbital – dumbbell shaped, 3 p

orbitals per sublevel.d orbital – 5 d orbitals per sublevel.f orbital – 7 f orbitals per sublevel.

Models of the AtomIn any orbital, there can be a maximum

of two electrons. The maximum number of electrons that

can occupy an energy level is given by the formula 2n2, where n is the # of the energy level.

1st energy level up to 2 electrons2nd energy level up to 8 electrons3rd energy level up to 18 electrons4th energy level up to 32 electrons

Quick ReviewMax of 2 electrons per orbital“s” sublevel – 1 orbital per sublevel (up

to 2 total electrons)“p” sublevel – 3 orbitals per sublevel

(up to 6 total electrons“d” sublevel – 5 orbitals per sublevel

(up to 10 total electrons)“f” sublevel – 7 orbitals per sublevel

(up to 14 total electrons)

Electron Arrangements in AtomsElectron configuration – the way in

which electrons are arranged in energy levels outside of the nucleus.

Orbital notation – a way of showing the electron configuration using arrows to represent each electrons and boxes to represent each orbital.

Electron Arrangements in Atoms RulesAufbau principle – electrons enter

orbitals of lowest energy first.

Pauli exclusion principle – an atomic orbital may hold at most two electrons.Electrons within the same orbital have

opposite spins.

Hund’s rule – one electron must be put in each orbital of a sublevel before any one orbital can have two electrons in it.

Orbital NotationsWhen writing orbital notations, use one

arrow to represent each electron.Electrons must enter the lowest energy

sublevel possible before moving to a higher energy sublevel

Even if you don’t have enough electrons to fill each orbital of a sublevel, you must still show that those orbitals exist.

The total number of arrows (electrons) must be equal to the atomic # for each element.

Types of Electrons in ArrangementsShared electrons – orbitals where there

are two electrons (arrows) with opposite spins.

Unshared electrons – when an orbital only has one electron in it.

Shared pair of electrons – any orbital that contains two electrons.

Orbital NotationsWrite the orbital notation for oxygen.

Write the orbital notation for aluminum.

Write out the orbital notation for cobalt

Electron Arrangements in AtomsElectron Configurations – the way in which

electrons are arranged around the nucleus of an atom. Each configuration has 3 parts:

1s2

“1” represents the energy level, “s” represents the sublevel, and “2” represents the number of electrons in that sublevel

The total of superscripts is equal to the atomic number for the element.

Electron Arrangements in Atoms1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p

Electron ConfigurationsWhich element is represented by the following

electron configuration:1s22s22p63s23p6

1s22s22p63s23p64s23d104p65s24d105p66s2

4f145d106p67s1

Electron ConfigurationsWrite the electron configuration for the

following elements:Sulfur

Gallium

Thorium

Platinum

Electron ConfigurationsWhat is wrong with each of the following

electron configurations?1s22s22p63s23p63d104s24p5

1s22s22p63s23p64s23d104p65s24d105p66s25d106p3

1s22s22p63s23p64s23d84p65s1

Noble Gas ConfigurationsNoble gas configurations are used as a

shorthand for long electron configurations.

Find the noble gas before the element you are writing the configuration for, put it in brackets, and then start with the next s sublevel to fill out the rest of the configuration.

Noble Gas ConfigurationsWrite the noble gas configuration for

the following elements:Sulfur

Iron

Thorium

Platinum

Noble Gas ConfigurationsWhat element is represented by the

following noble gas configuration:[Kr]5s24d105p2

[Ar]4s2

[Xe]6s24f145d6

Noble Gas ConfigurationsWhat is incorrect about the

following noble gas configurations?[Ar]2s22p2

[Kr]4d10

[At] 7s24f146d7

Electron ConfigurationElements in the same column on the

Periodic Table have Similar chemical and physical properties Similar valence shell electron configurationsSame numbers of valence electronsSame orbital typesDifferent energy levels

Valence electrons – outer energy level “s” and “p” sublevel electrons or electrons that are furthest away from the nucleus

Noble Gas Configurations & their relation to the Periodic TableLithium – [He]2s1

Sodium – [Ne]3s1

Potassium – [Ar]4s1

Rubidium – [Kr]5s1

Fluorine – [He]2s22p5

Chlorine – [Ne]3s23p5

Bromine-[Ar]4s23d104p5

Iodine-[Kr]5s24d105p5

s1

s2

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

p1 p2 p3 p4 p5 s2

p6

f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

1234567

Periodic TrendsAtomic radius – distance from the nucleus

of an atom to its valence electrons. The radius tells that size of the atom.

Moving from left to right across a period, atomic radius decreases.

Electrons within the same energy level don’t have as great of an effect on one another as electrons from different energy levels.

Trend in Atomic SizeIncreases down column

valence shell farther from nucleus

Decreases across period left to right

adding electrons to same valence shellvalence shell held closer because

more protons in nucleus

Periodic TrendsMoving down a group, atomic radius

increases.

The valence electrons get further and further from the nucleus because you are adding more energy levels. Therefore the radius of the atom increases.

Representation of Atomic Radii of the Main-Group Elements

Periodic TrendsExample: Put the following

elements in order of increasing atomic radius:

Zn, Sc, Se, K, Cs, O

Example: Put the following elements in order of decreasing atomic radius:

F, Cd, Ba, Ge, W, Cl

Periodic TrendsIonization energy – the energy

required to remove an electron from an atom (1st ionization energy).

Removing an electron creates a charge imbalance, so a cation (positive ion) is formed.

2nd Ionization energy – the energy required to remove two electrons from an atom.

Periodic TrendsMoving from left to right across a

period, ionization energy increases.

Within the same energy level electrons experience an increasing pull from the nucleus, so it takes more energy to remove them.

Periodic TrendsMoving down a group, ionization

energy decreases.

The valence electrons feel less and less pull from the nucleus as they get further from the nucleus.

Periodic Trends2nd ionization energy is always

greater than the 1st ionization energy.When you remove an electron from

an atom the number of protons becomes greater than the number of electrons. The remaining valence electrons move closer to the nucleus, making it harder to pull them off the atom.

Periodic TrendsAs electrons are removed, ionization

energy increases gradually until an energy level is empty, then it makes a big jump.

Pulling an electron off of a alkali metal (Group 1 elements) is easy. Trying to pull an electron off of a noble gas (Group 18 elements) takes much more energy.

Periodic TrendsExample: Put the following

elements in order of increasing ionization energy:

Sr, Cr, As, S, Rb, Cu

Example: Put the following elements in order of decreasing ionization energy:

O, V, K, P, Ga, Fr

Periodic Trends

Which of the following elements will have a very large second ionization energy? Third ionization energy?

Na, Al, Ne, Mg, Si

Periodic TrendsIonic radius – similar to atomic

radius but it is the radius for an ion instead of an atom.

Positive ions are always smaller than their neutral atoms, and negative ions are always larger than their neutral atoms.

As you go down a group, ionic radius increases.

Comparison of Atomic and Ionic Radii

Periodic TrendsAs you go from left to right across a

period, positive ions decrease in size.

Negative ions also decrease as you go across a period, but they start off being much larger than positive ions.

Periodic TrendsPut the following ions in order of

increasing ionic radius:Hint: If all of the ions have the same

number of electrons, than the one with the highest number of protons has the smallest radius.

Na+1, Al+3, N-3, F-1, O-2, Mg+2

Periodic TrendsElectronegativity – how strongly the

nucleus of an atom attracts the electrons of other atoms in a bond.

Nonmetals tend to gain electrons when they form bonds, and have higher electronegativities than metals, which tend to lose electrons, when they form bonds.

Periodic TrendsMoving from left to right across a

period, electronegativity increases.

Moving down a group, electronegativity decreases.

Electronegativities of the Elements

Periodic TrendsPut the following elements in order

of increasing electronegativity:Fe, Si, O, Ba, Ca, Cs

Put the following elements in order of decreasing electronegativity:

Se, F, Ag, Pt, Fr, Sb