UNIT 2: THE ATOM

Textbook Chapter 5 & 13Review Book Topic 1

Atomic Fundamentals All matter is composed of tiny fundamental

particles called atoms 

Atom – smallest particle of an element that retains the properties of that element

Examples of atomic size:

A pure copper coin the size of a penny contains 2.4 x 1022 atoms

Compared to the Earth’s population (6 x 109 people), there are about 4 x 1012 as many atoms in the coin as there are people on Earth

If you could line up 100,000,000 copper atoms side by side, they would produce a line on 1 cm long

Early Models of the Atom Democritus of Abdera 

4th Century B.C. – Greece

First to suggest the existence of atoms as invisible and indestructible particles

Components: Fire, Earth, Wind and Water 

The real nature of atoms was not established for more than 2000 years later

Robert Boyle (1600s)Identified gold and silver as being

elemental

Not made of Earth, fire, wind or water

Considered the “Father of Modern Chemistry”

John Dalton (1803)Studied the ratios in compounds

Set the groundwork for the current concept of the atom

Dalton’s Atomic Theory1. Atoms cannot be broken down

2. Atoms of the same element are identical

3. Each element’s atoms are different

4. Atoms of different elements can chemically or physically combine to form compounds

Structure of the Atom Most of Dalton’s atomic theory is

accepted today but atoms can be broken down into even smaller particles:

ProtonsNeutronsElectrons

Electrons – negatively charged particles

Discovered by J.J. Thomson in 1897

○ Experimented with a cathode tube

A cathode ray is composed of negatively charged particles

Negative electrical charges repels the rays, while positive charges are attracted

Refer to pages 109-110 in your textbook

○ Called his atomic model the “plum pudding model”

Robert A. Millikan (1868-1953)

○ An electron carries exactly one unit of negative charge

○ Mass of an electron is 1/1840 the mass of a hydrogen atom

Protons and Neutrons

Atoms have no net electric charge

Entire atom is neutral so:(+) charges = (-) charges

So # protons = # electrons in a neutral atom!

Protons – positively charged subatomic particles, each with a mass about 1840 times that of an electron

Neutron – subatomic particles with no charge but with a mass nearly equal to that of a proton

So if an atom is overall neutral….

# protons (+) = # electrons (-)

Properties of Subatomic Particles

Particle Symbol Relative electrical charge

Relative mass (mass of

proton = 1)

Actual mass

(g)

Location

Electron e- 1 - 1/1840 amu 9.11 x 10-28

Outside

Proton p+ 1 + 1 amu 1.67 x 10-24

Nucleus

Neutron n0 0 1 amu 1.67 x 10-24

Nucleus

Structure of the Atomic Nucleus Ernest Rutherford (1909)

Proposed an atom model where theelectrons surround a dense nucleus and all remaining areas are empty space

Directed a narrow beam of alpha (α) particles (with a + charge) at a very thin sheet of gold foil

A majority of alpha particles passed straight through the gold atoms, without deflecting while some particles bounced off the gold foil at very large angles or back toward the source

He found that:

○ An atom is mostly empty space (99.9% of the alpha particles went straight through the gold foil)

○ (+) charges and the mass of the atom are concentrated in the nucleus (0.1% of the alpha particles deflected)

See animation on his experiment

Nucleus – the central core of an atom and is composed of protons and neutrons (tiny when compared to the size of the atom overall - remember the Bill Nye video!)

Brain Pop

Niels Bohr (1913)

Proposed that electrons are arranged in circular paths, or orbits, around the nucleus with a fixed energy (no energy can be lost by the electron)

○ Model is patterned after the motions of planets around the sun (“planetary model”)

Erwin Schrödinger & Albert Einstein (1926)

Used a mathematical equation to describe the location and energy of an electron in a hydrogen atom

 Known as the “quantum mechanical model” or “wave

mechanical model”

○ Previous models were mostly physical based on the motion of large objects

 Restricts the energy of electrons to certain

values

Does not define an exact path an electron around the nucleus

Estimates the probability of finding an electron in a certain volume of space surrounding the nucleus (in a “fuzzy cloud”)

○ Cloud is more dense where there is high probability

○ Regions of probability are called “orbitals”

○ Energy level (of an electron) – region around the nucleus where the electron is likely to be moving

An electron can jump from one energy level to another

To move from one energy level to another, an electron must gain or lose just the right amount of energy

○ Quantum (of energy) – amount of energy required to move an electron from its present energy level to the next higher one

The higher the energy, the farther away from the nucleus the electron is located

Energy levels become more closely spaced the farther they are from the nucleus (stairs example)

History of Atomic Models

Subatomic Particles in Detail

Protons(+) charged particle found in the nucleus of an atom

# of protons in an atom determines which element it is

○ # protons = atomic number

○ Found on the periodic table of elements

If the # protons change in an atom, it becomes a different element

Neutrons

Subatomic particles with no charge

Number of neutrons can change within the atom without changing the element involved

Isotope – two atoms with the same number of protons but different numbers of neutrons

○ Ex. C-12 & C-14

○ Both are carbon so they have 6 protons...but C-14 has 2 more neutrons than C-12

○ Brain Pop

Neutrons Continued….

Mass Number = # protons + # neutrons

To find the # of neutrons:

○ Mass # - Atomic # = # Neutrons

○ (p+ + n0) - (p+) = (n0)

Symbol used to show atomic number and mass number of an element:

Can also be written as:

Element – (mass #)

Ex. Carbon-12 versus Carbon-14 (shows the change in mass number in these two isotopes)

Electrons(-) charged particles of an atom

To simplify the wave-mechanical model, we will draw “rings” around the nucleus to show electron configuration

Atoms overall are NEUTRAL (so p+ # = e- #) but….

○ Ions – atoms that have gained or lost electrons

Charge is determined by the different between # of protons and # of electrons

To Do: Complete the yellow table by using your

periodic table in your reference tables

You may work with a partner

The first row is done for you….

Isotopes Isotopes of an element have different #s of

neutrons thus also a different mass #

Yet….isotopes of the same element have identical chemical behaviors

Ex. Hydrogen isotopes:H-1H-2 (deuterium)H-3 (tritium)

Examples:Find the number of neutrons in an atom of Se-79

A neutral atom with 6 electrons and 8 neutrons is an isotope of….

Note: a few isotopes are listed on Table N in your reference tables…..

Atomic Mass The weighted average of the naturally

occurring isotopes of an element

Atomic mass unit (amu) – 1/12th the mass of a carbon-12 atom

In nature, most elements occur as a mixture of isotopes but one is more abundant than the others (which is the mass we estimate as a whole # from the periodic table)

To calculate atomic mass, you need:# of stable isotopes of the element

Mass of each isotope (in amu units)

Natural percent abundance of each isotope

Multiply atomic mass of each isotope by its abundance, expressed as a decimal, then add the results

Note:

Mass number is the whole number which is found by rounding the atomic mass on the periodic table for the element

Atomic mass is the average mass of all of the isotopes of the element

Example #1:Chlorine has two isotopes: chlorine-35 and

chlorine-37 (75% and 25%)

Which should the weighted average be closer to?.....Cl-35

To calculate:

Relative abundance MassCl-35 .75 x 35 = 26.25Cl-37 .25 x 37 = + 9.25

35.50

Example #2:Calculate the atomic mass of the two isotopes

of Boron: B-10 (19.78%) and B-11 (80.22%)

Relative abundance MassB-10 .1978 x 10 = 1.978B-11 .8022 x 11 = + 8.8242

10.8022

Example #3:

Atomic mass can be calculated from the mass and abundance of naturally occurring isotopes. Carbon has two naturally occurring stable isotopes. Most carbon atoms – 98.89% - are C-12, while the remaining 1.108% are C-13. What is the atomic mass of carbon?

Homework

Page 7 in REVIEW BOOK, questions #13-26- or-

Isotope Practice Worksheet

Ions Overall, atoms are neutral UNLESS they are

considered ions…

An ion is an atom with a (+) or (-) charge

(+) ions contain more protons (e- is lost)

(-) ions contain more electrons (e- is gained)

Examples:K+

Cl-

Mg2+

I-

Atomic Orbitals In the Bohr and quantum mechanical models, energy

levels of electron are designated by (n) – the principal quantum number

Each principal quantum number refers to a major energy level or orbital, represented by rings around the atom

Assigned in order of increasing energy (n = 1, 2, 3, 4) as distance from the nucleus increases

Atomic orbital – a region of space around the nucleus of an atom where there is a high probability of finding an electron

Letters are used to denote the shape:○ s orbital – spherical ○ p orbital – dumbbell-shaped○ d orbital – clover-leaf shapes○ f orbital – too complex to visualize

Nodes – in p and d orbitals, there are regions close to the nucleus where the probability of finding the electron is very low

Summary of Principal Energy Levels, Sublevels, and Orbitals

PrincipalEnergy Level

Number ofSublevels

Type ofSublevel

Shape ofSublevel

n = 1 1 S

n = 2 2 s, p

n = 3 3 s, p, d

n = 4 4 s, p, d, f

The maximum number of electrons that can occupy a principal energy level is given by the formula 2n2, where n is the principal quantum number

The number of electrons allowed in each of the first four energy levels are as follows:

Energy level n 1 2 3 4

Maximum number of electrons allowed

2 8 18 32

Increasing energy(Increasing distance from nucleus)

Homework Draw orbital diagrams for each element

with an atomic number between 1-20 on a blank sheet of paper

As (n) increases, all previous sublevels of lower orbitals must be included

Ground State – occurs when e- occupy the lowest available orbital

Electrons can gain or lose a specific amount of energy to move among atomic sublevels

Excited State – occurs when e- absorb energy and temporarily move to a higher energy level

Heat, light, electricity allow e- to move into increasing energy levels

The excited e- quickly returns to a lower energy level, emitting the same amount of energy it absorbed in the form of light (Infrared, ultraviolet or visible)

The light emitted is of multiple wavelengths (colors) and can be collected by a spectrometer to form a bright line spectra

○ Each bright line spectra is unique to a specific type of atom (can be used to identify elements!)

Examples of e- in excited states

Electron Configuration Distribution of electrons in an atom around

the nucleus

A complete electron configuration of an atom is shown by writing symbols for all the occupied sublevels in sequence

Shown on the periodic table for each element

Ex. Oxygen – 1s22s22p4 or 2-6

Can be shown in shorthand notation:Coefficient represents the principal energy level

(n = 1, 2, 3, etc.)

Sublevel shape is designed by s, p, d, or f

Superscript represents the # of e- in that sublevel

The sum of the superscripts equals the number of electrons total in the atom

Rules:

Electrons enter the orbitals of the lowest energies first

No more than two electrons can be placed into any orbital box

A single e- must be placed into each orbital box of a given sublevel before pairing takes place

Order of sublevels

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10…

Additional Info:s orbital can hold a max of 2 e-

p orbital can hold a max of 6 e-

d orbital can hold a max of 10 e-

f orbital can hold a max of 14 e-

So…… n =1 can hold just 2 e- since it contains an s orbital; n = 2 can hold 8 e- since is contains both an s and p orbital, n = 3 holds 18 e- because it has s, p, and d orbital…..etc.

Example: Helium

Example: Sulfur

Example: Nickel

Example: Calcium

Ex. Try these!

Carbon

Sodium

Silver

Ex. Try these!

Carbon - 1s2 2s2 2p2

Sodium - 1s2 2s2 2p6 3s1

Silver - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9

Using the electron configuration, you can determine whether an atom is in its ground or excited states

Indicated if any lower energy sublevel is not completely full yet the next sublevel contains that e-

Ex. Which of the following is in an excited state?1s22s22p2

1s22s22p1

1s22s22p53s2

1s22s22p63s1

Using the electron configuration, you can determine whether an atom is in its ground or excited states

Indicated if any lower energy sublevel is not completely full yet the next sublevel contains that e-

Ex. Which of the following is in an excited state?1s22s22p2

1s22s22p1

1s22s22p53s2

1s22s22p63s1

Homework

Review book topic 1, read pages 12-15

Complete electron configuration worksheet

Lewis Dot Structures Diagrams the show the valence

electrons of an element

Element symbol is used to represent the atom’s nucleus and all inner orbital electrons

Valance electrons are shown using small dots around the element’s symbol (to a max of 8 dots)

Valence electrons - # of e- on the last ring (highest energy level)

Can be found using the element’s electron configuration

Within a group (vertical columns) on the periodic table, each element has the same # of valence electrons

Determines the element’s chemical properties

Octet rule:

Atoms tend to achieve the electron configuration of a noble gas (8 valence e-)

Ions are formed to when an atom gains or loses electrons in order to reach this “octet”

○ Plays a role in ionic bonding (which we will learn about later)

Examples:

K Br Al P

F Si Sn At

Xe In Ca Bi

H He