1. Bettelheim, Brown, Campbell and Farrell Chapter 2
Atoms
Bettelheim, Brown, Campbell and Farrell
Chapter 2

2. Classification of Matter

3. Classification of Matter
Element: Pure substance made up of “identical” atoms
116 known elements
88 occur in nature; others man-made
Represented by one or two letter symbols
First letter CAPITALIZED; second letter small

4. Elements found in Nature
Monatomic elements:
Exist as single atoms
Diatomic elements:
Occur as diatomic molecules (pairs of atoms)
H2, N2, O2, F2, Cl2, Br2, and I2
Polyatomic elements:
Have three or more atoms per molecule
O3, P4, S8, diamond

5. Classification of Matter
Compound: Pure substance made up of two or more elements in a fixed ratio by mass
Formula of a compound: Gives the ratio of each element (uses atomic symbols).
Examples: NaCl H2O

6. Classification of Matter
Mixture: Combination of two or more pure substances
Substances may be present in any mass ratio
Each substance has a different set of physical properties
Each substance keeps its own properties and identity
Can separate mixture into the individual substances by using the physical properties of the individual substances in the mixture

7. Types of Mixtures
Heterogeneous mixture: Substances are not evenly distributed throughout
Homogeneous mixture (Solution): Substances are evenly distributed throughout.

8. Dalton’s Atomic Theory - 1805
All matter is composed of very tiny particles (atoms)
All atoms of the same element are the same (same chemical properties)
Compounds are formed by the chemical combination of two or more different kinds of atoms
A molecule is a tightly bound combination of two or more atoms that acts as a unit

9. Evidence for Dalton’s Theory
Law of Conservation of Mass
Mass can be neither created or destroyed
Law of Constant Composition
Compounds have a definite composition by mass

10. .

11. Subatomic Particles The unit of mass is the atomic mass unit (amu)
one amu is defined as one-twelfth the mass of an atom of carbon with 6 protons and 6 neutrons in its nucleus
1 amu = x g

12. A typical atom

13. Mass and Atomic Numbers
Mass number: Sum of the number of protons plus neutrons in the nucleus of an atom
Atomic number: Number of protons in the nucleus of an atom
Notation for single atom or isotope

14. Isotopes
Isotopes: atoms with the same number of protons but a different number of neutrons
Most elements found as mixtures of isotopes
Atomic weight: weighted average of masses of isotopes of an element

16. Classification of Elements
Metals
Solids (except Hg), shiny, conduct electricity, ductile, and malleable
Tend to give up electrons to form positive ions
Nonmetals
On right side of Periodic Table (except H)
Brittle, dull, poor conductors of electricity
Tend to accept electrons to form negative ions

17. Classification of Elements
Metalloids (also called semi-metals)
B, Si, Ge, As, Sb, Te
Have some properties of metals and some of nonmetals
Silicon is a semiconductor
Does not conduct electricity at low voltages, but becomes a conductor at higher voltages

18. Classification of Elements

19. Classification of Elements, part 2
Much info known about elements by mid 1800s
Tried to organize elements in logical way
Dmitri Mendeleev’s pattern worked best
Noted that certain properties tended to recur periodically
Took elements in order of increasing mass
Placed elements with similar properties in same column

20. Modern Periodic Table
Elements arranged in order of increasing atomic number (# of protons)
Elements with repeating properties are in the same group or family (column or vertical row)
Elements in same period (horizontal row) change properties as you go across the period

21. .

22. Periodic Properties
Mendeleev looked at both chemical and physical properties

23. Trend within Same Group
The alkali metals, Group 1A elements
Melting and boiling points decrease as you go down table

24. Group 1A: Alkali Metals
React with halogens to form compounds such as NaCl
Form +1 ions
Very reactive

25. Examples of Periodicity
Fluorine, chlorine, bromine, and iodine fall in the same column

26. Examples of Trends in Periodicity
The halogens, Group 7A elements
Melting and boiling points increase as you go down table

27. Group 7A: Halogens
Halogens exist as diatomic molecules, such as Cl2, F2, etc
Halogens react with group I metals compounds such as NaCl
Form -1 ions
Very reactive

28. Examples of Periodicity
The noble gases, Group 8A elements
Melting and boiling points increase as you go down table

29. Group 8A: Noble Gases Do not react with other elements
Do not form ions

30. Why do elements in group have similar properties?
Outermost Electron configuration is the same for all elements in a group
Electron configuration: the arrangement of electrons outside the nucleus

31. Electron Configuration
Electrons are distributed in shells about the nucleus

32. Electron Configuration
Each shell (principal energy level) has a different maximum number of electrons it can hold

33. Electron Configuration
Shells (principal energy levels) are subdivided into orbitals

34. Electron Configuration
Different kinds of orbitals have definite (and different) shapes and orientations in space

35. Electron Configuration
Each orbital has a different energy
Lowest energy fills first

36. Electron Configuration: Building Atoms
Rule 1: Lowest energy orbitals fill first
The first three energy level orbitals fill in the order 1s, 2s, 2p, 3s, and 3p
Rule 2: Each orbital can hold a maximum of two electrons (with opposite spins)
Rule 3: Orbitals of equal energy each add one electron first, then add second electron to fill them completely.

38. Order of Filling Orbitals
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s Don’t just fill up one shell and then the next shell completely Lowest energy fill first

39. Types of orbitals per shell

40. Orbital # orbitals/shell # electrons/ orbital type s 1 2 p 3 6 d 5 10 f 7 14

41. Electron Configuration
Orbital box diagrams
a box represents an orbital
an arrow represents an electron
a pair of arrows with heads in opposite directions represents a pair of electrons with paired spins
Example: carbon (atomic number 6)

42. Electron Configuration Notation
(Compact shorthand)
3p4
3 is principal energy level (shell)
p is type of orbital (subshell)
Superscript 4 show that there are four electrons in 3p orbitals

43. Electron Configuration
Noble gas notation
the symbol of the noble gas immediately preceding the particular atom indicates the electron configuration of all filled shells
Example: carbon (atomic number 6)

44. Write the electron configuration for Vanadium (# 23)

46. Shorthand way of remembering
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f

47. .

49. Electron Configuration
Lewis dot structures for the Group 1A (alkali) metals

50. Lewis Dot Structures
Atomic symbol represents nucleus and core electrons
Valence electrons shown as dots around symbol

51. Ionization Energy
Ionization energy: the energy required to remove the most loosely held electron from an atom in the gaseous state
example: when lithium loses one electron, it becomes a lithium ion; it still has three protons in its nucleus, but now only two electrons outside the nucleus

52. Ionization Energy
Ionization energy is a periodic property

53. Ionization Energy Ionization energy is a periodic property
Increases as you move from left to right
Increases as you move up the table

54. Electron Configuration
Valence shell: the outermost incomplete shell
Valence electron: an electron in the valence shell
Core electron: Electron inside the outermost shell

55. Electrons in Energy Levels
The energy of electrons in an atom is quantized
An electron can have only certain allowed energies
Ground state: the electron configuration of lowest energy
Excited state: Electron has more than lowest possible energy