1. Elements, Atoms & Ions Chapter 41

2. Elements Over 112 known, 88 of which occur naturally.
others are man-made (made in labs)
Abundance is the mass percent in the Earth’s crust, oceans and atmosphere
oxygen most abundant element and accounts for about 20% of Earth’s atmosphere
the abundance and form of an element varies from place to place but given as an overall average
Each element has a unique symbol
The symbol of an element may be one letter or two
if two letters in the symbol, only the first letter is capitalized. 2

3. The Periodic Table

6. Dalton’s Atomic Theory
1. Elements are composed of atoms
2. All atoms of a given element are identical
all carbon atoms have the same characteristics of every other atom of carbon
3. Atoms of a given element are different from those of any other element
carbon atoms have different characteristics than oxygen atoms 3

7. John Dalton (1766 – 1844) was an English scientist who made his living as a teacher in Manchester.

8. Dalton’s Atomic Theory
4. Atoms of one element combine with atoms of other elements to form compounds.
Law of Constant Composition
all samples of a compound have the same proportion of the elements as in any other sample of that compound
Chemical Formulas 4

9. Dalton pictured compounds as collections of atoms
Dalton pictured compounds as collections of atoms. Here NO, NO2, and N2O are represented.

10. Dalton’s Atomic Theory
5. Atoms are indivisible in a chemical process.
all atoms present at the beginning of a chemical process must also be present at the end of the process.
atoms are not created or destroyed, they must be conserved.
atoms of one element cannot be turned into atoms of another element
You cannot turn atoms of lead into atoms of gold 5

11. Formulas Describe Compounds
Compound - distinct substance that is composed of the atoms of two or more elements and always contains exactly the same relative masses of those elements.
Compounds are described by the elements in them and how many atoms of each element are in that compound.
Chemical formula – indicates the type and number of each element in a given compound.
if there is only one atom of an element, the number is not written after the symbol, it is an implied 1. 6

12. Atomic Structure History
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

13. Schematic of a cathode ray tube.

14. Conclusions from the Study of the Electron
Cathode rays have identical properties regardless of the element used to produce them.
Therefore, all elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

15. The Electron Tiny, negatively charged particle
Very light compared to the mass of the rest of the atom
1/1836th the mass of a proton 8

16. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model (easier to think of as “chocolate chips" in chocolate chip cookie dough.

17. Rutherford’s Gold Foil Experiment

18. Rutherford’s Gold Foil Experiment
bullet = alpha particles, target atoms = gold foil
 particles are positively charged
gold atoms are about 50 larger than a particles.
Alpha particles are helium nuclei
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

19. (a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct. (b) Actual results.

20. Rutherford’s Results Over 98% of the  particles went straight through
About 2% of the  particles went through but were deflected by large angles
About 0.01% of the  particles bounced off the gold foil
Most of the volume of the atom is empty space 11

21. Rutherford’s Nuclear Model
The atom contains a tiny dense center called the nucleus
the volume is about 1/10 trillionth the volume of the atom
The nucleus is essentially the entire mass of the atom (extremely dense)
The nucleus is positively charged
the amount of positive charge of the nucleus balances the negative charge of the electrons
The electrons move around in the empty space of the atom surrounding the nucleus 12

22. Structure of the Nucleus
The nucleus was found to be composed of two kinds of particles
Some of these particles are called protons
charge = +1
mass is about the same as a hydrogen atom
Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons
The other particle is called a neutron
has no charge
has a mass slightly more than a proton 13

23. The Modern Atom
We know atoms are composed of three main pieces - protons, neutrons and electrons
The nucleus contains protons and neutrons
The nucleus is only about cm in diameter
The electrons move outside the nucleus with an average distance of about 10-8 cm
therefore the radius of the atom is about 100,000 times larger than the radius of the nucleus 14

24. A nuclear atom viewed in cross section.

25. Atomic Particles Particle Charge Mass # Location Electron -1
Electron cloud
Proton +1 1
Nucleus
Neutron

27. Atomic Structure
The number of protons in an atom of a given element is the same as the atomic number (Z).
found on the Periodic Table, whole # for each element
Element
# of protons
Atomic # (Z)
Carbon 6
Phosphorus 15
Gold 79 15

28. Atomic Structure
Mass number = protons + neutrons; always a whole number.
# of Neutrons = mass number - # of protons
Atomic mass – larger number in each element’s box on the periodic table. If you round the atomic mass of an element to the closest whole number you generally get the mass # for that element.

29. Atomic Structure # of Electrons = # of protons if the atom is neutral
If the chemical symbol is written with a charge, representing an ion, the charge indicates the number of electrons that have been added or removed from the atom.

30. Atomic Structure
If the ion has a positive charge (cation), subtract that charge from the # of protons to get the number of electrons.
If the ion has a negative charge (anion), add that charge number to # of protons to get the number of electrons.
# of Electrons = # protons – charge
Charge = # protons - # electrons

31. Nuclear Symbols Mass number (p+ + no) Element symbol Atomic number
Charge (if any)
Element symbol
Atomic number
(number of p+)

32. Finding # of particles and charge
2713Al+3
3517Cl-1
13756Ba+2

33. Isotopes
Isotopes - atoms of an element with the same number of protons and electrons, but different numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

34. Two isotopes of sodium.

35. Isotopes Examples
3517Cl 3717Cl
H-1 H-2 H-3
Copper – 63 Copper – 65

36. Determining Average Atomic Mass
To determine the average atomic mass for an element, you must
Multiple the percentage (percent abundance) of each isotope of the element by its mass number.
Add the products of the multiplications together.
Divide by 100.
Your answer should be very close to the atomic mass of the element for that element

37. Average Atomic Mass Examples
Find the average atomic mass of each of the following elements from their percentages and mass numbers.
69.17% 63Cu and 30.83% 65Cu
5.85% Fe-54, 91.75% Fe-56, 2.12% Fe-57 and 0.28% Fe-58

38. Elements Arranged in a pattern called the Periodic Table
Elements listed by increasing atomic number.
Originally elements were listed by increasing atomic mass (Dmitri Mendeleev), but later it was found that by placing the elements in order of increasing atomic number (Moseley).
Position on the table allows us to predict properties of the element 16

39. Mendeleev’s Original Periodic Table

40. The modern periodic table.

41. The Modern Periodic Table
Elements with similar chemical and physical properties are in the same column
Columns are called Groups or Families
Rows are called Periods
The properties of the elements in one period are very similar to the element directly above or below it. 17

42. Elements Metals Nonmetals Metalloids (Si, Ge, As, Sb, Te, Po, At)
about 75% of all the elements
lustrous, malleable, ductile, conduct heat and electricity
Nonmetals
dull, brittle, insulators, do not conduct heat and electricity
Metalloids (Si, Ge, As, Sb, Te, Po, At)
also know as semi-metals
some properties of both metals & nonmetals

43. Figure 4.12: The elements classified as metals and as nonmetals.

44. Important Groups Noble Metals Ag, Au, Pt
Group 18 = Noble Gases
He, Ne, Ar, Kr, Xe, Rn
all colorless gases at room temperature
very non-reactive, practically inert
found in nature as a collection of separate atoms uncombined with other atoms
Noble Metals
Ag, Au, Pt
all solids at room temperature
least reactive metals
found in nature uncombined with other atoms 19

45. Platinum is a noble metal used in jewelry and in many industrial processes.

46. The Modern Periodic Table
Main Group = Representative Elements
“A” columns (Groups 1,2 and 13-18)
Transition Elements
“B” columns (Groups 3-12) all metals
Bottom rows = Inner Transition Elements = Rare Earth Elements (Lanthanides and Actinides)
metals
really belong in Period 6 & 7 18

47. Figure 4.11: The periodic table.

48. Important Groups - Halogens
Group 17 = Halogens
very reactive nonmetals
react with metals to form ionic compounds
Fluorine = F2
pale yellow gas
Chlorine = Cl2
pale green gas
Bromine = Br2
brown liquid that has lots of brown vapor over it
Only other liquid element at room temperature is the metal Hg
Iodine = I2
lustrous, purple solid 20

49. Other Important Groups
Alkali metals (Group 1) – extremely reactive metals; generally react with halogens to form “salts.” Form ions with a +1 charge.
Alkali earth metals (Group 2) – not as reactive as alkali metals. Form ions with a +2 charge.
Transition metals – less reactive than metals of Groups 1 & 2. Form ions that usually have more than one possible charge.

50. Natural State of Elements
Most elements are solids at room temp. and can exist as single atoms.
All metals except for mercury and gallium are solids at room temp.
Nonmetals can be found in all three common states of matter; sulfur (solid), bromine (liquid) and oxygen (gas).

51. Argon gas consists of a collection of separate argon atoms.

52. In solid metals, the spherical atoms are packed closely together.

53. (a) Sodium chloride (common table salt) can be decomposed to the elements (b) sodium metal (on the left) and chlorine gas.

54. Liquid bromine in a flask with bromine vapor.

55. Diatomic Molecules
Diatomic molecules – exist only as pairs of atoms of these particular elements. They cannot exist as single atoms.
Remember the “7” on the Periodic Table plus Hydrogen to remember all of the diatomic molecules.
H2, N2, O2, F2, Cl2, Br2, I2

56. Figure 4.11: The periodic table.

57. Gaseous nitrogen and oxygen contain diatomic (two-atom) molecules.

58. The decomposition of two water molecules (H2O) to form two hydrogen molecules (H2) and an oxygen molecule (O2).

59. Allotropes
Many solid nonmetallic elements can exist in different forms with different physical properties, these are called allotropes
the different physical properties arise from the different arrangements of the atoms in the solid
Allotropes of Carbon include
diamond
graphite
buckminsterfullerene 21

60. Graphite and diamond, two forms of carbon.

61. The three solid elemental forms of carbon (allotropes).