1. Standard 1 Atomic Structure
Chapters 4-6

2. Periodic Table. Non-metals Metals Nobel gases Alkaline earth metals
Alkali metals
halogens
Semi-metals
Metal/non-metal
boundary.
Transition metals

3. Summary 1
Which elements are semi-metals?

4. 1b: groups of the Periodic Table
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
Non-metals:
poor conductors
Mostly liquid/gas
gain electrons
Example = nitrogen
Halogens:
Extremely reactive
Gain 1 electron
Mostly gases
Example = fluorine

5. Summary 2 Describe the differences between metals and non-metals.
Give an example of a metal
Give an example of a non-metal

6. 1c: Periodic Groups Alkali metals Transition 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

7. Summary 3
Which group of metals are most reactive?

8. 1a: organization of the periodic table
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.

9. Summary 4
Name another element that would have similar chemical properties to chlorine.
Name an atom that is in the same period as chlorine.

10. C The Periodic Table: organizes elements according to atomic number
Atomic number = number of protons
Atomic number 6 C
12.011 1 2 3 4 5 6 7 8 9 10

11. 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)

12. Isotopes
ex:

13. Summary 5
What is the mass number for each isotope of neon shown in the example?
What is the atomic mass for neon?

14. 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

15. Summary 6
How many electrons are in a magnesium atom?

16. When an atom gains or loses electrons it becomes an ion
Ion = charged particle
number electrons ≠ number protons
symbol
symbol Na
Na+

17. Summary 7
If a magnesium atom loses two electrons, how many electrons will this magnesium ion have?

18. Valence electrons are: responsible for chemical behavior of atom
used for chemical bonding
located in the outer orbital
1 valence e valence e-

19. Summary 8 How many valence electrons does nitrogen have?
How many total electrons does nitrogen have?

20. 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

21. Summary 9
What causes an emission spectrum?

22. 1c: Periodic Trends
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.

23. General trend for electronegativity:
Increasing electronegativity
Increasing
Note: for noble gases electronegativity = zero

24. Summary 10 Which is more electronegative: iodine or chlorine?
Which is more electronegative: argon or chlorine?

25. 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.

26. General trend for ionization energy:
Increasing ionization energy
Increasing
Note: noble gases have a high ionization energy

27. Summary 11 Which has a higher ionization energy: iodine or chlorine?
Which has a higher ionization energy: argon or chlorine?
Which has a lower ionization energy: chlorine or magnesium?

28. General trend for atomic size (volume)
Decreasing atomic size
decreasing
Increasing

29. Summary 12 Which is larger: magnesium or calcium?
Which is larger: magnesium or chlorine?

30. General trend for ionic size.
When atoms lose electrons they get much smaller
When atoms gain electrons they get much larger

31. Summary 13
Why is Na+ smaller than Na?

32. Standard 1e: The structure of an atom
All the mass of an atom is in the nucleus (Protons & neutrons are in the nucleus)
In between the nucleus and the electrons there is only empty space

33. Summary 14
Which particles inside the atom have mass?

34. 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’
Rutherford’s gold-foil experiment yielded evidence of the atomic nucleus. a) Rutherford and his coworkers aimed a beam of alpha particles at a sheet of gold foil surrounded by a fluorescent screen. Most of the particles passed through the foil with no deflection at all. A few particles were greatly deflected. b) Rutherford concluded that most of the alpha particles pass through the gold foil because the atom is mostly empty space. The mass and positive charge are concentrated in a small region of the atom. Rutherford called this region the nucleus. Particles that approach the nucleus closely are greatly deflected.

35. Summary 15
How does Rutherford’s experiment demonstrate that an atom is mostly empty space?