Standard 1Atomic Structure

Chapters 4-6

Nobel gaseshalogensSemi-metalsTransition metals

Alkaline earth metalsAlkali metals

Metal/non-metalboundary.

Metals

Non-metalsPeriodic Table.

Summary 1

• Which elements are semi-metals?

• Metals:– Good conductors– Solid (except mercury)– Lose electrons– Example = aluminum

• Semi-metals (metalloids):– Have properties of both

metals and non-metals– Common use =

– semi-conductors

– Example = silicon

1b: groups of the Periodic Table

• Non-metals:– poor conductors– Mostly liquid/gas– gain electrons– Example =

nitrogen

• Halogens:– Extremely reactive– Gain 1 electron– Mostly gases– Example = fluorine

Summary 2

1. Describe the differences between metals and non-metals.

2. Give an example of a metal3. Give an example of a non-metal

1c: Periodic Groups• Alkali metals

– Extremely reactive

– Lose 1 electron– Example: sodium

• Alkaline earth metals– Reactive– Lose 2 electrons– Example: calcium

• Transition metals– Can lose different

numbers of electrons

– Example: copper

• Noble gases– Extremely un-

reactive– Gases!– Example: helium

Summary 3

• Which group of metals are most reactive?

• The Periodic Table: organizes elements in groups and periods.

• Groups/families: elements have the same physical and chemical properties.

• Rows/periods: elements have the same number of electron shells.

1a: organization of the periodic table

Summary 4

1. Name another element that would have similar chemical properties to chlorine.

2. Name an atom that is in the same period as chlorine.

• The Periodic Table: organizes elements according to atomic number

•Atomic number = number of protons

1

3 4 10

2

97 865

6

C12.011

Atomic number

Mass• Mass number: the number of protons

and neutrons in an atom (units = amu)• Atomic mass (shown on the periodic

table): the average mass of all isotopes • Isotope: an atom with the same number

of protons and a different number of neutrons

• Note: atomic mass generally increases across the periodic table but not always… (look at atomic number 27&28, 52&53)

Isotopesex:

Summary 5

1. What is the mass number for each isotope of neon shown in the example?

2. What is the atomic mass for neon?

Standard 1d: electrons• All atoms have an equal number of

protons and electrons– Atoms are electrically neutral

•Atoms have no charge•Symbol: Ne

An equal number of positive protons and negative electrons results in zero charge

Summary 6

• How many electrons are in a magnesium atom?

• When an atom gains or loses electrons it becomes an ion– Ion = charged particle

•number electrons ≠ number protons

Na Na+

symbol symbol

Summary 7

• If a magnesium atom loses two electrons, how many electrons will this magnesium ion have?

1 valence e- 4 valence e-

• Valence electrons are:• responsible for chemical behavior of atom • used for chemical bonding• located in the outer orbital

Summary 8

1. How many valence electrons does nitrogen have?

2. How many total electrons does nitrogen have?

Identifying Atoms by Emission Spectrum:•Adding energy ‘excites’ electrons.•Electrons release energy when they return to the ‘ground state’ (lowest energy level)•Released energy = ‘emission spectrum’ •Each atom has a unique emission spectrum•Scientists use this information in many ways:

•CSI can identify elements in an unknown sample •Astronomers can identify elements in stars across the universe

Summary 9

What causes an emission spectrum?

• Electronegativity: The ability of an atom to attract an electron

• Example: chlorine is very electronegative because it wants to ______ an electron.

• Example: sodium is not very electronegative because it wants to ______ an electron.

1c: Periodic Trends

• General trend for electronegativity:

Increasing electronegativity

Incre

asin

g

Note: for noble gases electronegativity = zero

Summary 10

1. Which is more electronegative: iodine or chlorine?

2. Which is more electronegative: argon or chlorine?

• Ionization energy: the energy needed to remove an electron from an atom

• Example: fluorine has a high ionization energy because it wants to ______ an electron.

• Example: potassium has a low ionization energy because it wants to ______ an electron.

• General trend for ionization energy:

Increasing ionization energy

Incre

asin

g

Note: noble gases have a high ionization energy

Summary 11

1. Which has a higher ionization energy: iodine or chlorine?

2. Which has a higher ionization energy: argon or chlorine?

3. Which has a lower ionization energy: chlorine or magnesium?

• General trend for atomic size (volume)

Decreasing atomic size

Incre

asin

gdecreasing

Summary 12

• Which is larger: magnesium or calcium?

• Which is larger: magnesium or chlorine?

General trend for ionic size.• When atoms lose electrons they get

much smaller • When atoms gain electrons they get

much larger

Summary 13

Why is Na+ smaller than Na?

1. All the mass of an atom is in the nucleus (Protons & neutrons are in the nucleus)

2. In between the nucleus and the electrons there is only empty space

Standard 1e: The structure of an atom

Summary 14

Which particles inside the atom have mass?

Earnest Rutherford

Rutherford demonstrated that the entire atom is 10,000 times larger than the nucleus• The rutherford experiment:• A stream of positive particles (alpha

particles) is aimed at a piece of gold foil.• Only 1 in 8000 particles is deflected (pass

close to the gold nucleus).• All other particles travel through ‘empty

space’

Summary 15

• How does Rutherford’s experiment demonstrate that an atom is mostly empty space?