Chapter 5 Atomic Structure Objective (1) Summarize Dalton’s atomic theory Objective (2) Describe the size of an atom Early Models of the Atom 1. Democritus.

Chapter 5 Atomic Structure

Objective (1) Summarize Dalton’s atomic theoryObjective (2) Describe the size of an atom

Early Models of the Atom 1. Democritus First suggested the existence of small

particles, called atoms

2. John Dalton a. Performed experiments to test his atomic theory

b. Studied the ratios in which elements

combine in chemical reactions

c. Formulated theories from the findings

d. Atomic Theory

1. All matter is composed of tiny indivisible particles called atoms.

2. All atoms of a given element are identical. Atoms of a specific element are different from any other element.

3. Elements can be chemically combined to form compounds.

• Just how small is an atom? 1. A Cu penny composed of pure copper 2. Grind the penny into fine powder, each

speck of powder is made smaller 3. The smallest piece left is an atom

4. The Cu penny contains 2.4 x1022 atoms 5. However, if you lined up 100,000,000

copper atom side by side it would only be 1 cm long

6. Atoms are very small

• Objectives:3. Distinguish among protons, electrons, and

neutrons in terms of relative mass and charge

4. Describe the structure of an atom, including the location of the protons, electrons, and neutrons with respect to the nucleus

Electrons 1. negatively charged subatomic particles

2. J.J. Thomson

- Discovered the electrons - His experiment: passed an electrical

current through gases at low pressure

• Cathode ray tubes

Results: 1. Regardless of the types of electrode,

the same particles appeared

2. The particles were repelled by the negative plate of a magnet.

3. The particles caused a paddle-wheel to spin

Conclusions:1. The particles were in all kinds of matter.2. The particles were negatively charged.3. The particles had mass.4. He proposed the plum pudding model

of matter.

• Protons and neutrons 1. Chadwick: discovered the neutron neutron: no charge subatomic particle

2. Rutherford: discovered the nucleus, and that it was positive

: performed the Gold-Foil Experiment

What he did: 1. Shot alpha particles (+ charge) at a sheet

of gold

2. Picture:

Results:1. Most of the alpha particles passed

straight through the gold 2. A few of the alpha particles were slightly

deflected 3. A very small number of alpha particles

hit and bounced back off the foil

Conclusion 1. He proposed the nuclear model of the

atom. 2. Matter is mostly empty space (with light

electrons) 3. Scattered throughout are small areas of

positive 4. Very dense matter called the nucleus.

mass of atoms comes from :

volume of atoms comes from:

Particle Symbol Relative Electrical charge

Actual mass (g)

Electron e- 1- 9.11 x 10-28

Proton p+ 1+ 1.67 x 10-24

Neutron N0 0 1.67 x 10-24

• Objectives: 5. Explain how the atomic number identifies an

element.

6. Use the atomic number and mass number of an element to find the number of

protons, electrons, and neutrons.

1. the number of protons in the nucleus of an atom

2. In a neutral atom it also equals the electrons

3. Ex. K atomic # = 19 p+ = 19 e_ = 19

Atomic number

Mass number 1. the total number of protons and neutrons in

the nucleus

2. mass number = p+ + n0

3. neutrons = mass number – p+

• How to write an atom:

1.

2. A Z

X Where A = mass # Z = atomic # X = any element

3. Hydrogen – 1 ( mass number)

4. Fill in the table:

Element Atomic number

Mass number

Protons Electron Neutron

15 16

89 39

K 30 19

• What is a neutral atom?

- It has the same number of p+ as e−

- Ex. Ca atomic # = 20

p+ = 20 so you have +20 e- = 20 so you have -20

What is the total charge? 0

Objectives:7. Explain how isotopes differ and why the

atomic masses of elements are not whole numbers.

8. Calculate the average atomic mass of an element from isotope data.

Isotopes 1. atoms with the same number of protons but different number of

neutrons

2. Identified by the mass

3. Example: Chlorine-35 Chlorine-37 mass number

Atomic mass unit– weighted average mass of the isotopes of that element

1. Ex. Cl 35.453 amu

75% is chlorine-35 25% is chlorine-37

2. (mass)(% of element) + (mass)(% of element) = (total mass)

* change % to a decimal

Examples:3. Ex. Element X has 2 natural isotopes. The

isotope with a mass of 10.012 amu (X-10) has a relative abundance of 19.91%. The isotope with a mass of 11.009 amu (X-11) has a relative abundance of 80.09%. Calculate the atomic mass of this element.

4. An element, E, has an atomic mass of 18.40 and consists of two isotopes: E-17 with a mass of 16.95 and E-20 with a mass of 19.35. How much E-20 does this element contain?

E-20 = xE-17 = (100-x)

(19.35)(x) + (16.95)(100-x) = (100)(18.40)

Ion

1. atoms lose or gains one or more electron

2. Mg+2 number of protons = 12 number of electrons = 10

charge of ion = +2 3. + ions : loses electrons 4. - ions : gains electrons

3. Examples Atomic # Mass # p n e

27 Al3+

13

70 Zn30

• Periodic Table 1. Mendeleev - created the first periodic table - used atomic mass to place elements - left blanks for undiscovered elements

2. Henry Moseley - determined the atomic number

• The modern periodic table

• Characteristics 1. Rows or periods : there are seven

2. Groups, families or columns: there are 18

3. Periodic law: there is a repeating pattern of physical and chemical

properties

4. Metals - have high electrical conductivity - high luster (shiny) -ductile - malleable - all solids at room temperature except: mercury (Hg)

5. Nonmetals - nonlustrous - poor conductors - brittle - can be solids, liquids, or gases

6. Metalloids - have both metal and nonmetal properties - touch the zigzag line