Atomic Structure Dr. Steve Badger. As we work through this… Find a Periodic Table of the Elements in the text and keep it handy for reference Use your.

Atomic Structure

Dr. Steve Badger

As we work through this…

• Find a Periodic Table of the Elements in the text and keep it handy for reference

• Use your copy of the handout entitled “Atomic Structure” to follow along and take notes

Consider a simple model

• The nucleus is in the center of the atom and contains – Positively charged protons (p+)and– Neutrons (n) with no charge

• Outside the nucleus, and at a great relative distance from the nucleus, we find negatively charged electrons (e—)

Atoms have a net charge of zero

• It is not correct to say, “atoms have no charge.” The protons and electrons that makeup the atom have a charge.

• Atoms can have a net charge of zero if and only if the number of electrons (—) equals the number of protons (+).

Atomic Number = #p+

• The atomic number is the number of protons.

• While it is true that the number of electrons in an atom is equal to the number of protons, it is not correct to say that the atomic number is the number of electrons.

• Look in the Periodic Table Find and find the atomic number of hydrogen, carbon, nitrogen, neon, calcium, and iron.

Protons

• The number of protons an atom contains will determine which element it is.

• Thus, an atom is nitrogen if and only if it has seven protons.

• Another way of saying this is, “an atom is nitrogen because it has seven protons.”

Electrons

• Electrons have a negative charge and are outside the nucleus

• Electronic structure determines reactivity

• That is, the number and arrangement of electrons—especially the outermost electrons—determines how an atom reacts

Electronic Structure

• The electronic structure determines whether an atom usually forms covalent bonds or ionic bonds

• Elements in the same family (column) typically react similarly

Three Subatomic Particles

• The mass of the proton and the mass of the neutron are essentially the same: 1 amu (atomic mass unit)

• Electrons have a mass so small (compared to that of the proton and neutron) that we can ignore their mass

• We say their mass approaches zero

Masses of subatomic particles

Particle Mass

Proton 1 amu

Neutron 1 amu

Electron —> 0 amu

The mass of an atom

• So the mass of an atom is essentially the sum of the masses of the protons and the masses of the neutrons

Consider these three atoms

1st atom 2nd atom 3rd atom

1 proton 1 proton 1 proton

1 electron 1 electron 1 electron

0 neutrons

1 neutrons

2 neutrons

What element is represented in the first column? What element is represented in the second column? What element is represented in the third column? How do you know the answer?

Why hydrogen?

• So all three of them stand for hydrogen

• Will all three atoms react the same?

• Yes. Why?

What’s different about these 3 atoms?

• The masses are different. • The 1st atom has a total mass of 1

amu• The 2nd atom has a mass of 2 amu• The 3rd atom has a mass of 3 amu

Atoms of the same element with different masses are called….?

Isotopes

• Isotopes have the same number of protons, but different number of neutrons—thus different masses

• Most elements exist in nature as two or more isotopes

• Are you familiar with any isotopes?Sure! Name one

Carbon-14

• Chemists write 14C but say “carbon-14”

• The 14 means that it has a total mass of 14 amu

• Since it is carbon, you know that it has six protons. (Verify this in the Periodic Table.)

Atomic Mass = #p+ + #n

• How many neutrons does 14C have?

• Think like this: The sum of the number of protons and the number of neutrons equals the atomic mass.

• Thus, 14 minus 6 equals 8…so 14C has 8 neutrons.

Bonding

• Most atoms form either ionic bonds or covalent bonds.

• A few atoms can form either type depending on what it is bonded to.

• And metals form “metallic bonds”

Ionic bonds

• Ionic bonds are formed when one atom loses an electron(s) and another atom gains an electron(s)

• Now the number of protons is not equal to the number of electrons, and this “atom” has a charge…

• …but we no longer call these atoms—they have a charge—so we label them ions

Ionic bonds

• Oppositely charged ions attract each other, and we call this an ionic bond

• Sodium chloride (NaCl) is an example of an ionic bond (Na+Cl—)

• Elements on the far left side of the Periodic Table typically form ionic bonds with elements on the far right side of the Periodic Table

Covalent bonds

• When two atoms share a pair of electrons, each atom behaves as if the pair of electrons belongs to it

• This shared pair of electrons is called a covalent bond

• The atoms that make up water (H2O) are covalently bonded to each other

Some of you may need to

• Read and read this handout• Spend extra time learning the

terms