An atom is the basic building What is an atom? All objects ...€¦ · • In 1805, English chemist John Dalton proposed and published he theory on atoms. • Dalton believed matter

• An atom is the basic building

block of matter.

• All objects are made of

atoms. The air, your desk and

all living things are made of

atoms.

• Atoms are EXTREMELY small.

One atom is only one ten-

billionth of a meter wide!

© Stephanie Elkowitz 1 Atoms & Reactions

What is an

atom?

• The idea of an atom was

first developed around 460

B.C. by Green philosopher

Democritus.

• Democritus believed that

you could not indefinitely

break an object in half. At

some point you get to the

smallest bit of matter, which

cannot be broken. He called

this bit of matter an atom.


• In 1805, English chemist John

Dalton proposed and

published he theory on atoms.

• Dalton believed matter was

made of extremely small

atoms. Atoms of the same

substance have identical size,

mass and other properties.

• Dalton also believed atoms

could not be created,

subdivided or destroyed but

could combine to form

different chemical substances. © Stephanie Elkowitz 3 Atoms & Reactions

• In 1897, English physicist J.J.

Thomson discovered the

electron.

• Thomson discovered that an

electron is a tiny, negatively

charged particle.

• From this finding, Thomson

proposed the first model of an

atom.


• J.J. Thomson came up with his

model of the atom in 1904.

• According to Thomson, an atom

was made of negatively charged

particles (electrons) embedded in a

“soup” of positive charges.

• This model suggested atoms

resembled plum pudding. The

electrons are “plums” surrounded

in a positively charged “pudding.”

For this reason, Thomson’s model

is called the plum pudding model.


• Ernest Rutherford proposed an

alternative model in 1911.

• According to Rutherford, an atom is

made of a central, positively

charged region. He called this

region the nucleus.

• Rutherford believed electrons

surrounded the nucleus as a

“cloud.”

• He also believed the nucleus of the

atom was small and dense

compared to the overall size of the

atom. © Stephanie Elkowitz 6 Atoms & Reactions

• Rutherford proved his model of the

atom with a famous experiment

known as the Gold Foil

Experiment.

• Rutherford worked with two other

scientists on this experiment: Han

Geiger and Ernest Marsden.

• In this experiment, Rutherford shot

a bean of positively charged

particles, called alpha particles, at

a thin piece of gold foil. He

recorded where the alpha particles

scattered as they struck the gold. © Stephanie Elkowitz 7 Atoms & Reactions

Rutherford, Geiger & Marsden

• If Thomson’s model was correct,

all the alpha particles should pass

straight through the foil as the

particles struck the gold.

• So, why would they pass straight

through? According to Thomson’s

plum pudding model, the charges

are symmetrically and evenly

distributed through an atom. This

is why the alpha particles would

pass straight through.


• Rutherford observed most particles

pass straight through. However, some

particles scattered and some particles

“bounced off” the gold foil.

• Rutherford believe his atomic model

explained these experimental results:

– Most particles passed through because an

atom is mostly empty space.

– Some particles scattered because they

were deflected by negatively charged

electrons.

– Some particles reflected back because

they bounced off the positively charged

nucleus. © Stephanie Elkowitz 9 Atoms & Reactions

• In 1914, Neils Bohr modified

Rutherford’s model of the atom.

• Like Rutherford, Bohr believed

an atom is made of a small,

positively charged nucleus

surrounded by negatively

charged electrons.

• However, Bohr believed the

electrons traveled around the

nucleus in circular orbits.

• This model is known as the

Rutherford-Bohr Atomic Model.


• Today we refer to the

Rutherford-Bohr Atomic

Model when studying atoms.

• According to this model,

there are two major regions

of an atom:

1. The center region of the

atom is called the nucleus.

2. The surrounding area around

the nucleus is mostly empty

space. Electrons orbit the

nucleus within this area.


Nucleus

• Atoms are VERY tiny! They are so small that we use special

units of measurement to describe the size and mass of

atoms and particles within atoms.

• We measure the mass of atoms in atomic mass units (u).

• One atomic mass unit equals 1.66 x 10-27 kilograms.

• We measure the size of atoms in picometers (pm).

• One picometer equals one trillionth of a meter. In other

words, 1 picometer equals 1.00 x 10-12 meters.

• An atom is approximately 100 pm, or 1 ten-billionth of a

meter wide.


• Subatomic particles are the

particles found within an

atom.

• There are three subatomic

particles:

– Proton (p)

– Neutron (n)

– Electron (e-)

We often use abbreviations to

denote the different subatomic

particles. The abbreviations are

listed next to the particles above.


Proton Neutron

Electron

• Protons are found in the

nucleus of an atom.

• A proton is positively

charged.

• One proton has a +1

charge.

• A proton has a mass of 1

atomic mass unit.


Proton +

+

• Neutrons are also found in

the nucleus of an atom.

• A neutron is a neutral

particle. On other words, a

neutron has NO charge.

• A neutron has a mass of 1

atomic mass unit


Neutron

• Electrons are found in the

space surrounding the

nucleus.

• An electron is negatively

charged. One electron has a

-1 charge.

• An electron has negligible

mass. In other words, the

mass of an electron is SO

small, it is insignificant. To

keep it simple, we say

electrons have no mass.


Electron

–

–

• Because the nucleus

contains positive and

neutral particles, the net

charge of the nucleus is

POSITIVE.

• A force known as the

nuclear force holds the

proton(s) and neutron(s)

together in the nucleus.


• The space around the

nucleus is NEGATIVELY

charged because it contains

negatively charged

electrons.

• Electrons stay in orbit

around the nucleus due to

electromagnetic force. This

force “holds” an atom

together.


• Although we cannot view the

arrangement of particles

within an atom, we have an

instrument that allows us to

see the surface of atoms.

• A scanning tunneling

microscope (STM) can view

the atoms of the surface of an

object.

• This image was taken with an

STM. It shows the surface of a

piece of gold. You can actually

see individual atoms!


• All atoms have the same arrangement of subatomic

particles. A change in the amount of subatomic particles

will change the type of atom.

• An element is a type of atom with the same number of

protons.

• We also use the term element to describe a substance that

is made of the same type of atom.

• There are more than 100 different types of elements.

• An element is abbreviated with a one or two letter symbol.


• Examples of elements and their symbols are listed below:


Element Symbol

Hydrogen H

Carbon C

Oxygen O

Nitrogen N

Sodium Na

Lithium Li

Element Symbol

Neon Ne

Helium He

Fluorine F

Chlorine Cl

Aluminum Al

Iron Fe

• Electrons constantly move

around the nucleus of an

atom.

• They orbit the nucleus in

specific orbitals or shells that

surround the nucleus.


+

• A shell is also called an energy level.

• A shell is called an energy level

because it is associated with a

certain amount of energy.

• The shell closest to the nucleus has

the lowest energy. Electrons in the

first shell have the least amount of

energy.

• As you move away from the nucleus,

the energy associated with a shell

increases. Electrons in the shell

furthest away from the nucleus have

the most energy. © Stephanie Elkowitz 23 Atoms & Reactions

INCREASING

ENERGY

• Each shell has a

maximum amount of

electrons it can hold.

– The 1st shell holds 2 e-.

– The 2nd shell holds 8 e-.

– The 3rd shell holds 18 e-.

– The 4th shell holds 32 e-.

• In large atoms, you can

find up to 7 shells.

• No shell can hold more

than 32 electrons.


+

2 e-

8 e-

18 e-

32 e-

• How electrons are

arranged in the shells of

an atom is called

electron configuration.

• Electrons “fill” or take up

space in each shell.

• In general, electrons fill

lower shells first

because lower shells

have less energy.

• Once a shell is full,

electrons begin to fill the

next higher shell. © Stephanie Elkowitz 25 Atoms & Reactions

Fill this shell first...

then this shell...

and finally this shell.

• To show electron configuration, we

draw a Bohr Diagram.

• To draw a Bohr Diagram:

1. Draw a circle to represent the

nucleus of the atom.

2. Write the element’s symbol, number

of protons (p) and number of

neutrons (n) inside the circle.

3. Draw rings around the circle to

represent electron shells. Each ring

represents a different energy level.

4. Draw electrons as dots in the rings.

Remember, each “ring” can only hold

so many electrons.


Example: Oxygen (O)

8 protons, 8 neutrons, 8 electrons

Example 1: Draw a Bohr diagram of Carbon (C) that has 6

protons, 6 neutrons and 6 electrons.


Remember...

2 e- fill the 1st

shell, 8 e- fill

the 2nd shell

and 18 e- fill

the 3rd shell.

Example 1: Draw a Bohr diagram of Carbon (C) that has 6 protons,

6 neutrons and 6 electrons.


2 electrons completely fill the 1st shell

4 electrons partially fill the 2nd shell

C

6 p

6 n

Example 2: Draw a Bohr diagram of Sodium (Na) that has 11 protons,



Example 2: Draw a Bohr diagram of Sodium (Na) that has 11 protons,



Na

11 p

12 n

• The electrons found in the

outermost orbital are called

valence electrons.

• For this reason, the

outermost shell is called the

valence shell.

• The number of valence

electrons determines many

chemical properties of an

element.


valence

electrons

• An atom cannot have more

than 8 valence electrons.

• An atom with 8 valence

electrons is said to have a

full outer shell.

• For example, neon (Ne) has

8 valence electrons.


• Helium (He) has a full

valence shell with only two

electrons.

• Helium only has one shell.

The maximum amount of

electrons held in the first

shell is 2 electrons. Since

the shell holds the

maximum amount of

electrons it can hold, helium

is said to have a full valence

shell.


• There are two important numbers associated with an

element that help you determine the number of protons,

neutrons and electrons in a neutral atom of that element:

1. Atomic Number

2. Atomic Mass


• The atomic number is the number of protons found in an

element.

• Atoms of the same element have the same number of

protons and thus, the same atomic number.

• Example: The atomic number of Helium is 2. Therefore, there

are two protons in an atom of Helium.


• Atomic mass is the mass of an atom. It is measured in atomic

mass units (u).

• Atomic mass equals the sum of protons and neutrons in an

atom.

• Remember...

– The mass of one proton is 1 u.

– The mass of one neutron is 1 u.

– The mass of an electron is negligible. In other words, an

electron has insignificant mass and it does not contribute

to the weight of an atom.


Example: Carbon has 6 protons and 6 neutrons. What is

carbon’s atomic mass?

Protons + Neutrons = Atomic Mass

6 + 6 = 12

Atomic Mass = 12 atomic mass units


• If you know an atom’s atomic number and the number of its

neutrons, you can calculate atomic mass.

• Adding an element’s atomic number (which is equivalent to

the number of protons) to the number of neutrons equals the

elements atomic mass.


• If you know an atom’s atomic mass and atomic number, you

can calculate the number of neutrons in an atom.

• To find the number of neutrons, subtract atomic number from

the atomic mass.


Example: Nitrogen has an atomic mass of 14. Its atomic number

is 7. How many neutrons are found in this atom?

Atomic Mass − Atomic Number = # of neutrons

14 − 7 = 7

Number of neutrons = 7


• The number of protons is the same for all atoms of a

specific element.

• A change in the number of protons will change the type of

atom - the element to which the atom belongs.

• The number of neutrons and electrons in an atom can vary

without changing the element to which the atom belongs.

• A change in the number of neutrons will change the form of

an atom. Different forms of atoms are called isotopes.

• The change in the number of electrons will change the

electric charge of an atom. Atoms with an electric charge

are called ions.


• An isotope is a variety of an element with a different number

of neutrons.

• The name of an isotope is the name of the element followed

by a dash (-) and the atomic mass of the isotope.

• Example: Carbon can have 6, 7 or 8 neutrons. Carbon with

6 neutrons is called Carbon-12. Carbon with 7 neutrons is

called Carbon-13. Carbon with 8 neutrons is called Carbon-

14.


• The average atomic mass is the average of all the naturally

occurring isotopes of an element.

• To find average atomic mass:

1. Identify the different isotopes of the element and the

atomic mass of each isotope.

2. Multiple the atomic mass of each isotope by its percent

abundance (in decimal form). Percent abundance is the

percent the element is found in the natural world.

3. Find the sum of these values. The sum is equal to the

average atomic mass.


Example: Chlorine

Step 1: Identify the different

isotopes of the element and the

atomic mass of each element.

Step 2: Multiple the atomic mass of

each isotope by its percent

abundance (in decimal form).

Step 3: Add these values together

to find average atomic mass.


Chlorine-35 exists 76.77% abundance

Chlorine-37 exists 24.23% abundance

.7577 x 35 = 26.5195

.2423 x 37 = 8.9651

26.5195 + 8.9651 = 35.4846 ≈ 35.48

• Average atomic mass gives you an idea as to what the most

common isotope of an element is.

• To find the most common isotope from the average atomic

mass, round the average atomic mass to the nearest whole

number. This “trick” works for most elements.

• For example, the average atomic mass of carbon is 12.011.

Because this value is closest to 12, we can assume that the

most common isotope of Carbon is Carbon-12.


• Most isotopes are stable. Stable isotopes have a “happy”

balance of protons and neutrons.

• Some isotopes are not stable. An unstable isotope is called a

radioactive isotope. A radioactive isotope is called so

because it emits radiation.

• Radiation is the release of energy in the form of waves or

subatomic particles. Radiation is dangerous because it can

harm the cells of living things. Specifically, it can alter

genetic material (DNA) in cells.


• A radioactive isotope releases energy in order to become

stable. This process is called radioactive decay.

• During radioactive decay, the atom emits radiation.

Radiation can be high-energy waves and/or subatomic

particles.

• As the atom undergoes radioactive decay, it becomes stable.

As a stable isotope, the atom no longer emits radiation.


• An element can gain or lose electrons to form an ion.

• An ion is an atom with an electric charge.

• Neutral atoms have the same amount of protons and

electrons. These atoms have zero net charge.

• An ion does not have the same amount of protons and

electrons. These atoms have a positive or negative change.


• When an atom loses an electron, it loses a negative charge.

Therefore, the atom has more protons and is positive.

• A positive ion is called a cation.


• When an atom gains an electron, it gains a negative charge.

Therefore, the atom has more electrons and is negative.

• A negative ion is called an anion.


• To determine the electric charge of an ion, you must know

the difference between protons and electrons.

• If there are more protons, the ion is positive A “+” is used to

denote a positive ion.

• If there are more electrons, the ion is negative a “−” is used

to denote a negative ion.

• Write the charge as a superscript to the right of the atom.

– If the ion has a +1 charge, simply write +.

– If the ion has a −1 charge, simply write −.


Example: A sodium atom has 11 protons and 10 electrons. How

do you denote the charge on this atom?


Example: A sodium atom has 11 protons and 10 electrons. How


1. 11 − 10 = 1 (the different between protons and electrons)

2. There are more protons, so the ion is positive (+).

3. Na+ is the abbreviation for a sodium ion with a +1 charge.


Example: An oxygen atom has 8 protons and 10 electrons. How



Example: An oxygen atom has 8 protons and 10 electrons. How


1. 10 − 8 = 2 (the different between protons and electrons)

2. There are more electrons, so the ion is negative (−).

3. O−2 is the abbreviation for an oxygen ion with a −2 charge.


• An atom wants to gain or

lose electrons in order to

have a complete valence

shell.

• Remember... A complete

valence electron shell holds

8 electrons except for the

first shell - the first shell can

hold a maximum of 2

electrons


Why do

atoms

become

ions?

• An atom that has only one or two valence electrons tends to

lose those electrons and become a positive ion.

– Example: Lithium (Li) is an atom with 1 valence electron. Lithium

tends to lose this electron to become a +1 ion with a complete

valence electron shell.


• An atom that has nearly a full valence shell tends to gain

electrons and become a negative ion.

– Example: Fluorine (F) has 7 valence electrons. It tends to gain one

electron to become a −1 ion with a complete valence electron shell.


• There are more than 100 different elements. All the known

elements are organized in a table. This table is called the

Periodic Table of Elements.


• The periodic table displays

each element in a box.

• In most period tables, each

box gives 4 pieces of

information about an element:

①The name of the element

② The chemical symbol of the

element

③ The atomic number

④ The average atomic mass


1

H Hydrogen

1.00794

①

④

②

③

• Elements are presented in the periodic table by order of

increasing atomic number.


• The elements are arranged in rows and columns.

• The rows are called periods. The columns are called groups.


• Electron configuration

dictates how elements are

organized in the rows and

columns of the period table.

In general:

– Elements with the same

number of valence electrons

are placed in the same column.

– Elements with the same

number of shells are placed in

the same row.

• This organization explains the

shape of the periodic table.


Why is the

periodic table

shaped the

way it is?

• There are four major types of elements in the periodic table:

1. Metals 3. Metalloids

2. Nonmetals 4. Noble Gases


• Most elements are metals.

• Metals are hard, shiny (lustrous), malleable and ductile.

• Metal are good conductors and usually have a high density

and melting point.

• Examples: Sodium, Calcium, Iron, Gold


• Nonmetals make up the majority of the universe.

• Nonmetals are brittle and non-elastic. Most are dull.

• Nonmetals are poor conductors and usually have a low

melting point.

• Examples: Carbon, Phosphorus, Sulfur, Iodine, Chlorine

• Note: Hydrogen is a nonmetal even though it is located on

the periodic table where most of the metals are found.


• Metalloids have characteristics in-between metals and

nonmetals. They are shiny and have a “metallic

appearance.” Although they look like metals, most

metalloids are brittle, not malleable and fair conductors.

• There are 6 commonly recognized metalloids: Boron,

Silicon, Germanium, Arsenic, Antimony and Tellurium.

• Astatine and/or Polonium are sometimes included. The

periodic table used in this presentation includes Polonium.


• Noble gases are inert

(unreactive) substances.

• As their name suggests, Noble

gases are gases at room

temperature.

• Noble gases are poor

conductors of heat and

electricity.

• There are 6 Noble gases:

Helium, Neon, Argon, Krypton,

Xenon and Radon.


• Recall: Elements in the same group have the same number

of valence electrons.


• The number of valence electrons is easy to determine for

groups 1, 2 and 13 through 18.


1 V

ALE

NC

E e

-

2 V

ALE

NC

E e

-

3 V

ALE

NC

E e

-

4 V

ALE

NC

E e

-

5 V

ALE

NC

E e

-

6 V

ALE

NC

E e

-

7 V

ALE

NC

E e

-

8 V

ALE

NC

E e

-

• There are exceptions to this pattern in the center block of

elements and to the rows of elements below the table. These

elements are called transitional metals.

• Transitional metals have complicated electron configuration.

Electrons are NOT always added to the outermost orbital.


TRANSITIONAL METALS

TRANSITIONAL METALS

• Some groups have specific names because they share

similar properties:

– Group 1: Alkali Metals

– Group 2: Alkaline (Earth) Metals

– Group 3-12: Transitional Metals

– Group 17: Halogens

– Group 18: Noble Gases

• Elements in these groups share similar properties because

they have the same number of valence electrons.


• The named groups are outlined in the period table below.


HA

LO

GE

NS

NO

BLE

GA

SE

S

ALK

ALI M

ETA

LS

ALK

ALIN

E M

ETA

LS

TRANSITIONAL METALS

TRANSITIONAL METALS

• Alkali Metals are soft and shiny

metals found in Group 1.

• All Alkali Metals have one valence

electron. They readily lose this

electron to form a +1 ion.

• Hydrogen is NOT an Alkali Metal - it

is a nonmetal. It is placed in this

group because it has one valence

electron.

• Alkali Metals are extremely

reactive. If you put any of these

pure elements in water, they can

cause a huge explosion. © Stephanie Elkowitz 74 Atoms & Reactions

• Alkaline Earth Metals are Group 2

elements.

• These elements have two valence

electrons and readily lose these

electrons to form a +2 ion.

• Alkaline Earth Metals are the 2nd

most reactive elements.

• Within this group, you will find

Calcium. Calcium is important to

building your bones.


• Transitional Metals are metallic elements found in Groups 3-

12 and in the rows below the periodic table as well.

• Transitional metals have complicated electron configurations.

They do not always completely fill a shell before beginning to

fill the next shell.


• Halogens are nonmetals found in

Group 17.

• All of these elements have 7

valence electrons. They readily

gain one electron to form a -1 ion.

• Halogens are very reactive.

Fluorine is the most reactive

element in this group. The

reactivity of the elements in this

group decreases as you move

down the column.


• Noble Gases are elements found

in Group 18.

• All of these elements are gases.

• Noble Gases have a full outer

shell. Except for Helium, all the

elements in this group have 8

valence electrons. Helium has a

full first shell with 2 valence

electrons.

• Noble Gases are very unreactive.

They rarely react with other

elements.


• Below the periodic table are two rows of elements. These

rows make up the F Block of the periodic table.

• Elements in the F Block are part of the periodic table but it’s

easier to display them on the bottom of the table.

• All of the elements are transitional metals.

• Each row is considered a series of elements.

– The first row is called the lanthanide series.

– The second row is called the actinide series.


• The Lanthanide Series includes chemical elements with

atomic numbers 57 through 71.

• All the elements are shiny, silvery metals.

• These elements are rare.

• Lanthanide elements are useful to superconductors, glass

production and lasers.


• The Actinide Series includes chemical elements with atomic

numbers 89 through 103.

• All the elements are radioactive.

• Some of the actinide elements are synthetic - they are not

naturally found in Earth. Elements with atomic numbers 95

through 103 are synthetic elements.

• Plutonium is used in atomic weapons.

• Uranium has been used in atomic weapons. It is most often

used to produce electricity via nuclear power.


• There are other synthetic elements in the periodic table.

• In total, there are 24 synthetic elements. These elements

have atomic numbers 95 through 118.

• Synthetic elements do not occur naturally on Earth.

• All synthetic elements were made in a laboratory.

• All synthetic elements are unstable and radioactive. They

decay rapidly, some in only a few hundred microseconds.

• Some elements, such as Technetium and Plutonium, are

synthetically made. However, they are not purely synthetic.

They exist naturally in very small quantities on Earth.


• There are patterns to the properties of elements in the periodic table.

These patterns are called trends.


• Atomic radius is a measure of the size of an

atom.

• Specifically, atomic radius is a measure of

how “wide” an atom is from the center of its

nucleus to the outermost electron shell.


Atomic radius

• As you move down a column, atoms have more orbitals.

Orbitals increase the size of an atom. Therefore, atomic

radius increases as you move down a column.

• As you move across a row, atomic radius decreases. As an

atom gains electrons in the same orbital, the electrons are

more attracted to and “pulled” towards the positive nucleus.


• Atoms with nearly a full valence electron shell want to gain

electrons to completely fill the shell. These atoms have a

high electronegativity. Electronegativity is the tendency of an

atom to attract an electron. The greater the electronegativity,

of an atom, the mire likely the atom will attract an electron.


• As you move across a row, electronegativity increases. This

makes sense because atoms have more valence electrons.

• As you move down a column, an element’s electronegativity

decreases. It decreases because larger atoms have a more

difficult time attracting electrons.


• Ionization energy is the energy required to remove an electron

from an atom. The greater the ionization energy of an atom,

the “harder” it is to remove an electron from the atom.

• Atoms with 1 or 2 valence electrons easily give up electrons

while atoms with a nearly a full valence electron shell do not.


• As you move across a row, ionization energy increases. Atoms

have more valence electrons and do not want to give them up.

• As you move down a column, ionization energy decreases.

The size of an atom increases so it’s easier for an atom to

lose electrons.


• Elements, except hydrogen, on the left side of the period table

are metals. These elements have metallic character.

• Metallic character refers to the characteristics of metals, such

as luster, malleability, ductility and conductivity.

• Metalloids have some metallic character. Nonmetals do not

have metallic character, which is why they are called nonmetal.


• As you move across the periodic table (left to right),

elements lose metallic character. In other words, metallic

character decreases.

• Elements transition from metals to metalloids to nonmetals

as you move across the table. Metals have the most

metallic character. Nonmetals have no metallic character.


Documents

An atom is the basic building What is an atom? All objects ...€¦ · • In 1805, English chemist John Dalton proposed and published he theory on atoms. • Dalton believed matter