Unit 3- Atomic Structure

• Atom• Atomic mass• Atomic mass unit

(amu)• Atomic number• Electron• Excited state• Ground state

• Isotope • Law of definite

proportions• Mass number• Neutron• Orbital• Proton• Valence• Wave-mechanical

model

• Atomic theory- atoms are the building blocks of matter

• Law of definite proportions- a chemical compound will always have the same elements in the same proportions– Ex: glucose always C6H12O6

• Law of conservation of mass- mass cannot be created or destroyed in ordinary physical or chemical changes– Mass of products=mass of reactants– 89g oxygen + 11g hydrogen = 100g water

Changing views of atomic structure

1st- John Dalton (1803)5 principles:

1- all matter is composed of atoms that can’t be created, destroyed or divided

• now, atoms can be divided• technology allows us to create and destroy

2- atoms of an element are identical physically and chemically3- atoms of different elements differ in physical and chemical properties4- atoms of different elements combine into compounds (in whole number ratios)5- in chemical reactions atoms are combined, separated and rearranged, NEVER destroyed, created or changed

2nd- JJ Thompson 1897• Plum pudding model-• Experiment with cathode ray tube• Cathode ray (-) was deflected by magnetic field• Particles must have negative charge- discovered electrons• Electrons- particles with negative charge, outside the

nucleus

3rd- Ernest Rutherford 1909• Gold foil experiment- discovered positive nucleus and

that an atom is mostly empty space

• Nucleus- dense, central part of atom, positively charged

4th- Bohr (1913)• Electrons travel around the nucleus in specific

energy levels

Now- Wave-mechanical model• Determined that electrons have properties of mass

and also waves

• Difference from Bohr; instead of electrons being in fixed orbits they are in orbitals– Region in which an electron of a particular amount of

energy is most likely located (high probability)

The Nucleus• Consists of:

– Protons- positively charged• Makes up the atomic number of an atom• A neutral atom must have an equal number of

electrons and protons• Mass: 1amu (atomic mass unit)

– An atomic mass unit is 1/12th the mass of a C-12 atom

– Neutrons- neutral in charge• Helps hold protons together• Neutrons + protons = atomic mass of an atom

• Atomic number- # of protons as atom has– Also shows the number of electrons

• Ex: oxygen has an atomic number of 8; so 8 protons and 8 electrons

• Mass number- # of protons and neutrons in nucleus– Calculate neutrons by subtracting atomic number

from mass number

16 8 O

Isotopes

• atoms of the same element that have a different number of neutrons– Atomic number is the same– Atomic mass is different

***different elements can have the same mass number but can’t have

the same atomic number***

Calculating average atomic mass• Most elements are mixtures of isotopes• On periodic table- mass is the average

– If you know the percentage of each isotope you can calculate the average

– Ex: calculate average mass of nitrogen if a typical sample contains 99.63 % nitrogen-14 and 0.37 % nitrogen-15

• Step 1- convert percent to decimal99.63% = .9963 0.37% = .0037• Step 2- multiply decimals by masses.9963 x 14 = 13.95.0037 x 15 = .056• Step 3- add products together13.95 + .056 = 14.006

How many atoms make up that mass?

• Moles are used by scientists to make sense of the vast number of atoms in a sample– The amount of a substance– 6.022 x 1023 atoms/particles - Avogadro’s number– The number of atoms in exactly 12 grams of C-12– Ex: a dozen eggs, not 12 eggs

Converting with moles • Molar mass (gram formula mass) -the mass (g) in 1

mole of a substance– Ex: Carbon = 12g/mol Sulfur = 32 g/mol

• What is the mass in grams of 2.34 mol of lead?2.34mol x (207.2g/mol) = 484.85g Pb

• How many moles are in 222g of copper? 222g x (1mol/63.55g) = 3.49 mol Cu

Converting with molecules

• How many atoms are in .632 moles of boron?.632 moles x (6.022x1023) = 3.81x1023 B atoms• ** How many atoms are in 6g of neon?1st- how many moles? 6g x (1mol/20.18g) =.297mol Ne2nd- moles to atoms:.297mol x (6.022x1023atoms/1mol) = 1.79x1023

Ne atoms

Electron configuration

• Bohr - Electrons travel around the nucleus in specific energy levels– Closer to the nucleus- the lower the E level

• Like rungs on a ladder, the higher up (farther out) the more potential energy you have, can’t also be in between rungs

Background on light• Light is energy• Using a prism, light can be broken into different

colors.• Different colors of light have different amounts of

energy.• A full spectrum of light has many different

amounts of energy.

Flame tests

• Electrons are normally in a ground state (lowest PE), – if it gains energy it moves to an excited state (high PE)

• In it’s excited state it quickly falls back to ground state and releases ENERGY in the process

• This release of energy emits a certain wavelength of light

• Measuring an electrons potential energy tells how far from the nucleus it is

– When solutions containing metal ions are heated, they impart characteristic colors to a flame when the electrons fall back down to ground state

Sodium

Calcium

Barium

Potassium

Placing the Electrons

• The Aufbau Principle − electrons fill orbitals starting at the lowest available energy state before filling higher states.

• Hund’s Rule − unoccupied orbitals of a given energy will be filled before occupied orbitals of the same energy are reused. (like seats on a bus)

• The Pauli Exclusion Principle − no two electrons in an atom can occupy the same quantum state. (be in the same orbital and have the same spin)o An orbital can hold a maximum of two electrons

The location of electrons follows certain rules or principles:

Patterns• The first energy level has only one sublevel, s; the second energy level has

two sublevels, s and p; the third energy level has three sublevels, s, p, and, d; and so on.

• There is 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, and so on.• There are a maximum of two electrons per orbital.• Using this information, can you fill in the table below?

Principal Energy Level

(n)

Number of

Orbitals ( )

Electrons per Sublevel Maximum Number of Electrons

( )s p d f

1 3

Electrons in Each Locati

on

1 1 2 − − − − − − 2

2 4 2 6 − − − − − 8

3

4

5

6

7

g h i

5 7 9 11 13

916

n22n2

2536

49

1832

(50)

(72)

(98)

−2 6 10 − − −2 6 10 14 − − −

2 6 10 14 18 − −

2 6 10 14 18 22 −

2 6 10 14 18 22 26

An Analogy

• Each floor is like an energy level. (1-7)• Each apartment is like a sublevel. (s, p, d, or f)

o Higher energy levels are larger, so they have more sublevelsThe first level (floor) has one sublevel (apartment) − sThe second has two − s and pAnd so on

• Each bedroom is like an orbital. o “s” apartments have one bedroomo “p” apartments have three bedroomso “d” apartments have five bedrooms, ando “f” apartments have seven bedrooms

Sometimes it is helpful to think of the atom as an apartment house.

Expanding the Analogy• Imagine electrons are moving into the apartment complex

pictured below:o Electrons don’t like to waste energy climbing to apartments on

higher floors.o Electrons don’t like to waste energy caring for larger

apartments.o Electrons move into the most energy efficient apartments first.

• In what order do the apartments fill up?1

2 34 5

67

89

1011

12

Order of Filling• The electrons are arranged according to the following rules:

o the number of electrons equals the number of protons (atomic number)

o electrons occupy orbitalsin sequence beginning with those of the lowest energy.

o in a given sublevel, a second electron is not added to an orbital until each orbital in the sublevel contains one electron.

• One consequence of the orderof filling is that an outer shell never has more than eight electrons.

1s2

2s2 2p6

3s2 3p6 3d10

4s2 4p6 4d10 4f14

5s2 5p6 5d10 5f14

6s2 6p6 6d10

7s2 7p6

A Sample Electron Configuration• Consider the element scandium shown below:

• What do the rules for the order of filling show about the electron configuration?o First 1s fills with 2 electrons leaving 19o Second 2s fills with 2 electrons leaving 17o Next 2p fills with 6 electrons followed by

3s with 2 electrons leaving 9o Then 3p fills with 6 electrons followed by 4s

with 2 electrons leaving 1o Finally the remaining electron goes into 3d

• This gives an electron configuration of 1s22s22p63s23p64s23d1.

44.9559 +3

Sc212-8-9-2

1s

2s 2p

3s 3p 3d

4s

Types of Electron Configurations

• Sublevel Notation: sublevel notation shows how many electrons are in each sublevel

• Bohr Notation: Bohr notation shows the number of electrons in each shell or energy level

• Orbital Notation: Orbital notation shows the electrons and their spin in each orbital

1s22s22p63s23p64s23d11s22s22p63s23p64s23d1

2 − 8 − 9 − 2

1s___↑↓

2s___↑↓ ___↑↓

2p___↑↓___↑↓

3s___↑↓ ___↑↓

3p___↑↓___↑↓

4s___↑↓ ___↑ ___

3d ___ ___ ___

More on Orbital Notation

• Step 1: Determine the sublevel notation of irono 1s has room for 2 electrons, leaving 24.o 2s has room for 2 electrons, leaving 22.o 2p has room for 6 electrons and 3s has room for 2

more, leaving 14.o 3p has room for 6 electrons and 4s has room for 2

more, leaving 6.o 3d has room for the remaining 6o The electron configuration is 1s22s22p63s23p64s23d6.

• Step 2: Use the information from the sublevel notation to draw the orbital notationo Draw a horizontal line to represent each orbital in a sublevel, and label it.o Add one electron at a time to each orbital, represented by an up or down

arrow, in the same order as in the sublevel notation. Electrons in an orbital must have opposite spins (arrows in opposite directions). Follow Hund’s rule. Do not begin pairing electrons in an orbital until all the orbitals

in a sublevel have an electron

Drawing the Orbital Notation for Iron (Atomic Number = 26)

1s___

2s___ ___

2p___ ___

3s___ ___

3p___ ___

4s___ ___ ___

3d ___ ___ ___↑↓

1s

2s 2p

3s 3p 3d

4s

↑↓ ↑ ↑ ↑↓ ↓ ↓ ↑↓ ↑ ↑ ↑↓ ↓ ↓ ↑↓ ↑ ↑ ↑ ↑ ↑↓

Summary

• Draw the sublevel notation by following the order of filling.

• Draw the Bohr notation byadding together all the electrons in the same energylevel.

• Draw the orbital notation by taking the information from the sublevel notation.

1s2

2s2 2p6

3s2 3p6 3d10

4s2 4p6 4d10 4f14

5s2 5p6 5d10 5f14

6s2 6p6 6d10

7s2 7p6

Fe: 1s22s22p63s23p64s23d6

Fe: 2-8-14-2

1s___↑↓

2s___↑↓ ___↑↓

2p___↑↓___↑↓

3s___↑↓ ___↑↓

3p___↑↓___↑↓

4s___↑↓ ___↑ ___

3d ___ ___ ___↑ ↑ ↑ ↑↓Fe: