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## PowerPoint Slideshow about 'Atomic Structure' - PamelaLan

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Democritus

- Circa 400 BC
- Greek philosopher
- Suggested that all matter is composed of tiny, indivisible particles, called atoms

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Dalton’s Atomic Theory (1808)

- All matter is made of tiny indivisible particles called atoms.
- Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
- Atoms of different elements can combine with one another in simple whole number ratios to form compounds.
- Chemical reactions occur when atoms are separated, joined, or rearranged;however, atoms of one element are not changed into atoms of another by a chemical reaction.

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J.J. Thomson (1897)

- Determined the charge to mass ratio for electrons
- Applied electric and magnetic fields to cathode rays
- “Plum pudding” model of the atom

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Rutherford’s Gold Foil Experiment (1910)

- Alpha particles (positively charged helium ions) from a radioactive source was directed toward a very thin gold foil.
- A fluorescent screen was placed behind the Au foil to detect the scattering of alpha () particles.

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Rutherford’s Gold Foil Experiment (Observations)

- Most of the -particles passed through the foil.
- Many of the -particles deflected at various angles.
- Surprisingly, a few particles were deflected back from the Au foil.

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Rutherford’s Gold Foil Experiment (Conclusions)

- Rutherford concluded that most of the mass of an atom is concentrated in a core, called the atomic nucleus.
- The nucleus is positively charged.
- Most of the volume of the atom is empty space.

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Shortfalls of Rutherford’s Model

- Did not explain where the atom’s negatively charged electrons are located in the space surrounding its positively charged nucleus.
- We know oppositely charged particles attract each other
- What prevents the negative electrons from being drawn into the positive nucleus?

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Bohr Model (1913)

- Niels Bohr (1885-1962), Danish scientist working with Rutherford
- Proposed that electrons must have enough energy to keep them in constant motion around the nucleus
- Analogous to the motion of the planets orbiting the sun

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Planetary Model

- The planets are attracted to the sun by gravitational force, they move with enough energy to remain in stable orbits around the sun.
- Electrons have energy of motion that enables them to overcome the attraction for the positive nucleus

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Think about satellites….

- We launch a satellite into space with enough energy to orbit the earth
- The amount of energy it is given, determines how high it will orbit
- We use energy from a rocket to boost our satellite, what energy do we give electrons to boost them?

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Electronic Structure of Atom

- Waves-particle duality
- Photoelectric effect
- Planck’s constant
- Bohr model
- de Broglie equation

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Radiant Energy

- Radiation the emission of energy in various forms
- A.K.A. Electromagnetic Radiation
- Radiant Energy travels in the form of waves that have both electrical and magnetic impulses

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Electromagnetic Radiation radiation that consists of wave-like electric and magnetic fields in space, including light, microwaves, radio signals, and x-rays

- Electromagnetic waves can travel through empty space, at the speed of light (c=3.00x108m/s) or about 300million m/s!!!

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Waves

Waves transfer energy from one place to another

Think about the damage done by waves during strong hurricanes.

Think about placing a tennis ball in your bath tub, if you create waves at one it, that energy is transferred to the ball at the other = bobbing

Electromagnetic waves have the same characteristics as other waves

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Wave Characteristics

Wavelength, (lambda) distance between

successive points

2m

10m

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Wave Characteristics

- Frequency, (nu) the number of complete wave cycles to pass a given point per unit of time; Cycles per second

t=5

t=0

t=0

t=5

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Units for Frequency

- 1/s
- s-1
- hertz, Hz
- Because all electromagnetic waves travel at the speed of light, wavelength is determined by frequency
- Low frequency = long wavelengths
- High frequency = short wavelengths

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Electromagnetic Spectrum

- Radio & TV, microwaves, UV, infrared, visible light = all are examples of electromagnetic radiation (and radiant energy)
- Electromagnetic spectrum: entire range of electromagnetic radiation

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Electromagnetic Spectrum

Frequency Hz

1024 1020

1018 1016 1014

1012 1010 108 106

Gamma Xrays UV

Microwaves FM AM

IR

10-16 10-9 10-8 10-6 10-3 100 102

Wavelength m

Visible Light

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Notes

- Higher-frequency electromagnetic waves have higher energy than lower-frequency electromagnetic waves
- All forms of electromagnetic energy interact with matter, and the ability of these different waves to penetrate matter is a measure of the energy of the waves

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What is your favorite radio station?

- Radio stations are identified by their frequency in MHz.
- We know all electromagnetic radiation(which includes radio waves) travel at the speed of light.
- What is the wavelength of your favorite station?

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Velocity of a Wave

- Velocity of a wave (m/s) = wavelength (m) x frequency (1/s)
- c =
- c= speed of light = 3.00x108 m/s
- My favorite radio station is 105.9 Jamming Oldies!!!
- What is the wavelength of this FM station?

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Wavelength of FM

- c =
- c= speed of light = 3.00x108 m/s
- = 105.9MHz or 1.059x108Hz
- = c/ =3.00x108 m/s = 2.83m

1.059x1081/s

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What does the electromagnetic spectrum have to do with electrons?

- It’s all related to energy – energy of motion(of electrons) and energy of light

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States of Electrons

- When current is passed through a gas at a low pressure, the potential energy (energy due to position) of some of the gas atoms increases.
- Ground State: the lowest energy state of an atom
- Excited State: a state in which the atom has a higher potential energy than it had in its ground state

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Neon Signs

- When an excited atom returns to its ground state it gives off the energy it gained in the form of electromagnetic radiation!
- The glow of neon signs,is an example of this process

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White Light

- White light is composed of all of the colors of the spectrum = ROY G BIV
- When white light is passed through a prism, the light is separated into a spectrum, of all the colors
- What are rainbows?

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Line-emission Spectrum

- When an electric current is passed through a vacuum tube containing H2 gas at low pressure, and emission of a pinkish glow is observed.
- What do you think happens when that pink glow is passed through a prism?

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Hydrogen’s Emission Spectrum

- The pink light consisted of just a few specific frequencies, not the whole range of colors as with white light
- Scientists had expected to see a continuous range of frequencies of electromagnetic radiation, because the hydrogen atoms were excited by whatever amount of energy was added to them.
- Lead to a new theory of the atom

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Bohr’s Model of Hydrogen Atom

- Hydrogen did not produce a continuous spectrum
- New model was needed:
- Electrons can circle the nucleus only in allowed paths or orbits
- When an e- is in one of these orbits, the atom has a fixed, definite energy
- e- and hydrogen atom are in its lowest energy state when it is in the orbit closest to the nucleus

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Bohr Model Continued…

- Orbits are separated by empty space, where e- cannot exist
- Energy of e- increases as it moves to orbits farther and farther from the nucleus

(Similar to a person climbing a ladder)

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Bohr Model and Hydrogen Spectrum

- While in orbit, e- can neither gain or lose energy
- But, e- can gain energy equal to the difference between higher and lower orbitals, and therefore move to the higher orbital (Absorption)
- When e- falls from higher state to lower state, energy is emitted (Emission)

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Bohr’s Calculations

- Based on the wavelengths of hydrogen’s line-emission spectrum, Bohr calculated the energies that an e- would have in the allowed energy levels for the hydrogen atom

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Photoelectric Effect

- An observed phenomenon, early 1900s
- When light was shone on a metal, electrons were emitted from that metal
- Light was known to be a form of energy, capable of knocking loose an electron from a metal
- Therefore, light of any frequency could supply enough energy to eject an electron.

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Photoelectric Effect pg. 93

- Light strikes the surface of a metal (cathode), and e- are ejected.
- These ejected e- move from the cathode to the anode, and current flows in the cell.
- A minimum frequency of light is used. If the frequency is above the minimum and the intensity of the light is increased, more e- are ejected.

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Photoelectric Effect

- Observed: For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum, no matter how long the light was shone
- Why does the light have to be of a minimum frequency?

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Explanation….

- Max Planck studied the emission of light by hot objects
- Proposed: objects emit energy in small, specific amounts = quanta

(Differs from wave theory which would say objects emit electromagnetic radiation continuously)

Quantum: is the minimum quantity of energy that can be lost or gained by an atom.

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Planck’s Equation

- E radiation = Planck’s constant x frequency of radiation
- E = h
- h = Planck’s constant = 6.626 x 10-34 J•s
- When an object emits radiation, there must be a minimum quantity of energy that can be emitted at any given time.

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Einstein Expands Planck’s Theory

- Theorized that electromagnetic radiation had a dual wave-particle nature!
- Behaves like waves and particles
- Think of light as particles that each carry one quantum of energy = photons

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Photons

- Photons: a particle of electromagnetic radiation having zero mass and carrying a quantum of energy
- Ephoton = h

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Back to Photoelectric Effect

- Einstein concluded:
- Electromagnetic radiation is absorbed by matter only in whole numbers of photons
- In order for an e- to be ejected, the e- must be struck by a single photon with minimum frequency

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Example of Planck’s Equation

- CD players use lasers that emit red light with a of 685 nm. Calculate the energy of one photon.
- Different metals require different minimum frequencies to exhibit photoelectric effect

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Answer

- Ephoton = h
- h = Planck’s constant = 6.626 x 10-34 J•s
- c =
- c= speed of light = 3.00x108 m/s
- = (3.00x108 m/s)/(6.85x10-7m)
- =4.37x10141/s
- Ephoton= (6.626 x 10-34 J•s)(4.37x10141/s)

Ephoton= 2.90 x 10-19J

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Wave Nature of Electrons

- We know electrons behave as particles
- In 1925, Louis de Broglie suggested that electrons might also display wave properties

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de Broglie’s Equation

- A free e- of mass (m) moving with a velocity (v) should have an associated wavelength: = h/mv
- Linked particle properties (m and v) with a wave property ()

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Example of de Broglie’s Equation

- Calculate the wavelength associated with an e- of mass 9.109x10-28 g traveling at 40.0% the speed of light.
- 1 J = 1 kg m2/s2

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Answer

- C=(3.00x108m/s)(.40)=1.2x108m/s
- = h/mv
- = (6.626 x 10-34 J•s) =6.06x10-12m

(9.11x10-31kg)(1.2x108m/s)

Remember 1J = 1(kg)(m)2/s2

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Wave-Particle Duality

- de Broglie’s experiments suggested that e- has wave-like properties.
- Thomson’s experiments suggested that e- has particle-like properties
- measured charge-to-mass ratio

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Where are the e- in the atom?

- e- have a dual wave-particle nature
- If e- act like waves and particles at the same time, where are they in the atom?
- First consider a theory by German theoretical physicist, Werner Heisenberg.

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Heisenberg’s Idea

- e- are detected by their interactions with photons
- Photons have about the same energy as e-
- Any attempt to locate a specific e- with a photon knocks the e- off its course
- ALWAYS a basic uncertainty in trying to locate an e-

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Heisenberg’s Uncertainty Principle

- Impossible to determine both the position and the momentum of an e- in an atom simultaneously with great certainty.

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SchrÖdinger’s Wave Equation

- An equation that treated electrons in atoms as waves
- Only waves of specific energies, and therefore frequencies, provided solutions to the equation
- Quantization of e- energies was a natural outcome

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SchrÖdinger’s Wave Equation

- Solutions are known as wave functions
- Wave functions give ONLY the probability of finding and e- at a given place around the nucleus
- e- not in neat orbits, but exist in regions called orbitals

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SchrÖdinger’s Wave Equation

- Here is the equation
- Don’t memorize this or write it down
- It is a differential equation, and we need calculus to solve it

-h (ә2 Ψ )+ (ә2Ψ )+( ә2Ψ ) +Vψ =Eψ

8(π)2m (әx2) (әy2) (әz2 )

Scary???

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Definitions

- Probability likelihood
- Orbital wave function; region in space where the probability of finding an electron is high
- SchrÖdinger’s Wave Equation states that orbitals have quantized energies
- But there are other characteristics to describe orbitals besides energy

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Quantum Numbers

- Definition: specify the properties of atomic orbitals and the properties of electrons in orbitals
- There are four quantum numbers
- The first three are results from SchrÖdinger’s Wave Equation

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Quantum Numbers

- Principal Quantum Number, n
- Values of n = 1,2,3,…
- Positive integers only!
- Indicates the main energy level occupied by the electron

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Quantum Numbers

- Principal Quantum Number, n
- Values of n = 1,2,3,…
- Describes the energy level, orbital size

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Quantum Numbers

- Principal Quantum Number, n
- Values of n = 1,2,3,…
- Describes the energy level, orbital size
- As n increases, orbital size increases.

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Principle Quantum Number

- More than one e- can have the same n value
- These e- are said to be in the same e- shell
- The total number of orbitals that exist in a given shell = n2

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Quantum Numbers

- Angular momentum quantum number, l
- Values of l = n-1, 0
- Describes the orbital shape

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Quantum Numbers

- Angular momentum quantum number, l
- Values of l = n-1, 0
- Describes the orbital shape
- Indicates the number of sublevel (subshells)

(except for the 1st main energy level, orbitals of different shapes are known as sublevels or subshells)

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Orbital Shapes

- For a specific main energy level, the number of orbital shapes possible is equal to n.
- Values of l = n-1, 0
- Ex. Orbital which n=2, can have one of two shapes corresponding to l = 0 or l=1
- Depending on its value of l, an orbital is assigned a letter.

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Orbital Shapes

- Angular magnetic quantum number, l
- If l = 0, then the orbital is labeled s.
- s is spherical.

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Orbital Shapes

- If l = 2, the orbital is labeled d.
- “double dumbbell” or four-leaf clover

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Energy Level and Orbitals

- n=1, only s orbitals
- n=2, s and p orbitals
- n=3, s, p, and d orbitals
- n=4, s,p,d and f orbitals

Remember: l = n-1

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Atomic Orbitals

- Atomic Orbitals are designated by the principal quantum number followed by letter of their subshell
- Ex. 1s = s orbital in 1st main energy level
- Ex. 4d = d sublevel in 4th main energy level

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Quantum Numbers

- Magnetic Quantum Number, ml
- Values of ml = +l…0…-l
- Describes the orientation of the orbital
- Atomic orbitals can have the same shape but different orientations

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Magnetic Quantum Number

- s orbitals are spherical, only one orientation, so m=0
- p orbitals, 3-D orientation, so m= -1, 0 or 1 (x, y, z)
- d orbitals, 5 orientations, m= -2,-1, 0, 1 or 2

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Quantum Numbers

- Electron Spin Quantum Number,ms
- Values of ms = +1/2 or –1/2
- e- spin in only 1 or 2 directions
- A single orbital can hold a maximum of 2 e-, which must have opposite spins

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Electron Configurations

- Electron Configurations: arragenment of e- in an atom
- There is a distinct electron configuration for each atom

There are 3 rules to writing electron configurations:

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Pauli Exclusion Principle

- No 2 e- in an atom can have the same set of four quantum numbers (n, l, ml, ms ). Therefore, no atomic orbital can contain more than 2 e-.

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Aufbau Principle

- Aufbau Principle: an e- occupies the lowest energy orbital that can receive it.
- Aufbau order:

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Hund’s Rule

- Hund’s Rule: orbitals of equal energy are each occupied by one e- before any orbital is occupied by a second e-, and all e- in singly occupied orbitals must have the same spin

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Electron Configuration

- The total of the superscripts must equal the atomic number (number of electrons) of that atom.
- The last symbol listed is the symbol for the differentiating electron.

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Differentiating Electron

- The differentiating electron is the electron that is added which makes the configuration different from that of the preceding element.
- The “last” electron.
- H 1s1

He 1s2

Li 1s2, 2s1

Be 1s2, 2s2

B 1s2, 2s2, 2p1

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Orbital Diagrams

- These diagrams are based on the electron configuration.
- In orbital diagrams:
- Each orbital (the space in an atom that will hold a pair of electrons) is shown.
- The opposite spins of the electron pair is indicated.

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Orbital Diagram Rules

- Represent each electron by an arrow
- The direction of the arrow represents the electron spin
- Draw an up arrow to show the first electron in each orbital.
- Hund’s Rule: Distribute the electrons among the orbitals within sublevels so as to give the most unshared pairs.
- Put one electron in each orbital of a sublevel before the second electron appears.
- Half filled sublevels are more stable than partially full sublevels.

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Orbital Diagram Examples

- H _

1s

- Li _

1s 2s

- B __ __

1s 2s 2p

- N _

1s 2s 2p

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Dot Diagram of Valence Electrons

- When two atom collide, and a reaction takes place, only the outer electrons interact.
- These outer electrons are referred to as the valence electrons.
- Because of the overlaying of the sublevels in the larger atoms, there are never more than eight valence electrons.

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Rules for Dot Diagrams

:Xy:

. .

Px orbital

S sublevel electrons

. .

Py orbital

Pz orbital

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Rules for Dot Diagrams

- Remember: the maximum number of valence electrons is 8.
- Only s and p sublevel electrons will ever be valence electrons.
- Put the dots that represent the s and p electrons around the symbol.
- Use the same rule (Hund’s rule) as you fill the designated orbitals.

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Summary

- Both dot diagrams and orbital diagrams will be use full to use when we begin our study of atomic bonding.
- We have been dealing with valence electrons since our initial studies of the ions.
- The number of valence electrons can be determined by reading the column number.
- Al = 3 valence electrons
- Br = 7 valence electrons
- All transitions metals have 2 valence electrons.

V.Montgomery & R.Smith

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