Chapter 5 Atomic Structure Objective 1 Summarize Daltons

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 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 x 1022 atoms 5. However, if you lined up 100, 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: 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.

Particle Symbol Electron Proton Neutron ep+ N 0 Relative Electrical charge 11+ 0 mass of atoms comes from : volume of atoms comes from: Actual mass (g) 9. 11 x 10 -28 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.

Atomic number 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

Mass number 1. the total number of protons and neutrons in the nucleus 2. mass number = p+ + n 0 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 Mass Protons Electron Neutron number 15 16 89 K 30 39 19

• What is a neutral atom? - It has the same number of p+ as e− - Ex. Ca atomic # = 20 p+ = 20 e- = 20 so you have +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 = x E-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 # 27 Al 3+ 13 70 Zn 30 p n e

• 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