Atomic Structure Timeline From this part of the lecture, you will be expected to… draw the atomic models match scientists to their experiments and discoveries place the models in chronological order

Democritus (400 B.C.) Proposed that matter was composed of tiny indivisible particles Greek: atomos=small, solid, indestructible particles of different shapes and sizes Not based on experimental data

Alchemy (next 2000 years) Mixture of science and mysticism. Lab procedures were developed, but alchemists did not perform controlled experiments like true scientists.

John Dalton (1807) British Schoolteacher based his theory on others’ experimental data Billiard Ball Model atom is a uniform, solid sphere

History of Atomic Models 5 Main Points of Dalton’s Atomic Theory Elements are made of tiny, indivisible particles called atoms All atoms of a given element are identical Atoms of a given element are different than those of any other element Atoms of one element can combine with atoms of other elements to give compounds Atoms cannot be created, divided into smaller particles, nor destroyed in the chemical process; reactions change how atoms are grouped.

Problems with Dalton’s Atomic Theory? 1. Matter is composed, indivisible particles called atoms Atoms Can Be Divided, but only in a nuclear reaction 2. All atoms of a particular element are identical Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)! 3. Different elements have different atoms YES! 4. Atoms combine in certain whole-number ratios YES! Called the Law of Definite Proportions 5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements. Yes, except for nuclear reactions that can change atoms of one element to a different element

Henri Becquerel (1896) Discovered radioactivity spontaneous emission of radiation from the nucleus Three types: alpha () - positive beta () - negative gamma () - neutral

J. J. Thomson (1903) Cathode Ray Tube Experiments beam of negative particles Discovered Electrons negative particles within the atom Plum-pudding Model

J. J. Thomson (1903) Plum-pudding Model positive sphere (pudding) with negative electrons (plums) dispersed throughout

Ernest Rutherford (1911) Gold Foil Experiment Discovered the nucleus dense, positive charge in the center of the atom Nuclear Model

Ernest Rutherford (1911) Nuclear Model dense, positive nucleus surrounded by negative electrons

Niels Bohr (1913) Bright-Line Spectrum tried to explain presence of specific colors in hydrogen’s spectrum Energy Levels electrons can only exist in specific energy states Planetary Model

Niels Bohr (1913) Planetary Model Bright-line spectrum Planetary Model electrons move in circular orbits within specific energy levels

Bohr’s Model The electron must gain energy to move to a higher level and must lose energy to move to a lower level Bohr said electrons act more like waves on a string than particles Think of walking from floor to floor of the school building!

Erwin Schrödinger (1926) Quantum mechanics electrons can only exist in specified energy states Electron cloud model orbital: region around the nucleus where e- are likely to be found

Electron Cloud Model (orbital) Erwin Schrödinger (1926) Electron Cloud Model (orbital) dots represent probability of finding an e- not actual electrons

James Chadwick (1932) Discovered neutrons neutral particles in the nucleus of an atom Joliot-Curie Experiments based his theory on their experimental evidence

revision of Rutherford’s Nuclear Model James Chadwick (1932) Neutron Model revision of Rutherford’s Nuclear Model

Atomic Structure I. Structure of the Atom Chemical Symbols Subatomic Particles Atomic # and Atomic Mass Isotopes Average Atomic Mass Mass of Compounds

Metal that forms bright blue solid compounds. Chemical Symbols Capitals matter! Element symbols contain ONE capital letter followed by lowercase letter(s) if necessary. Metal that forms bright blue solid compounds. Co vs. CO Poisonous gas.

Atomic Structure An atom is the smallest unit of matter made of subatomic particles: protons, neutrons, and electrons

Subatomic Particles in a neutral atom Most of the atom’s mass. NUCLEUS ELECTRONS in a neutral atom PROTONS NEUTRONS NEGATIVE CHARGE POSITIVE CHARGE NEUTRAL CHARGE Most of the atom’s mass. Atomic Number equals the # of...

Electrons Bohr compared electrons to planets, saying electrons orbited the nucleus in specified and circular paths However, an electrons exact location cannot be determined

Orbitals Electrons exist in energy levels called orbitals The number of filled orbitals depends on how many electrons an atom has Electrons occupy the orbitals that have the lowest energy (closest to the nucleus) Four different kinds: s, p, d, f

Orbital Region where there is 90% probability of finding an electron. Can’t pinpoint the location of an electron. Density of dots represents degree of probability.

Orbtials s orbitals are spherical and can hold 2 electrons p orbtials are dumbell shaped and can hold six electrons

Orbital Orbitals have different shapes.

Valence Electrons Electrons located in the outer most orbital are called valence electrons These electrons determine the atom’s chemical properties and its abilities to form chemical bonds Atoms with the same # of valence electrons have similar properties

Bohr Model Diagrams Simplified energy levels using Bohr’s idea of circular orbits. Can replace with: 3p 4n Lithium Atomic #: 3 Mass: 7 # of p: 3 # of e: 3 # of n: 4 e- e- p n Maximum e- Level 1 2e- Level 2 8e- Level 3 18e- Level 4 32e- e-

Bohr Model Activity Choose a number between 1 & 18. Find your element by the atomic number you picked. Draw a Bohr Model diagram for your element on your marker board. Round off the mass listed on the table and subtract the atomic # to find the # of neutrons. Abbreviate the # of ‘p’ and ‘n’ in the nucleus. Have a partner check your drawing. Repeat with a new element.

III. Masses of Atoms Atomic Mass Mass Number Isotopes Ch. 10 - Atomic Structure III. Masses of Atoms Atomic Mass Mass Number Isotopes

Atomic Number Equals # of protons Equals # of electrons in a NEUTRAL atom Always a whole number

Atomic Mass atomic mass unit (amu) 1 amu = 1/12 the mass of a 12C atom 1 proton = 1 amu 1 neutron = 1 amu 1 electron=1/1822 amu Lightest subatomic particle is the electron 1 amu = 1.67  10-24 g © Addison-Wesley Publishing Company, Inc.

Mass Number Sum of the protons and neutrons in the nucleus of an atom. © Addison-Wesley Publishing Company, Inc. Always a whole number.

Mass Number # of protons + neutrons in atomic mass units (amu) Isotopes - atoms of the same element with different masses differ in number of neutrons Examples: Carbon-13 & Carbon-14, Boron-10 & Boron-11 Element-mass#

Calculating # of neutrons Mass # - # Protons Example: Aluminum 13 protons 27-13 = 14 neutrons

Isotopes Mass # Atomic # Atoms of the same element with different mass numbers. Nuclear symbol: Mass # Atomic # Hyphen notation: carbon-12

Isotopes © Addison-Wesley Publishing Company, Inc.

Isotopes Chlorine-37 atomic #: mass #: # of protons: # of electrons: # of neutrons: 17 37 20

Relative Atomic Mass 12C atom = 1.992 × 10-23 g atomic mass unit (amu) 1 amu = 1/12 the mass of a 12C atom 1 p = 1.007276 amu 1 n = 1.008665 amu 1 e- = 0.0005486 amu © Addison-Wesley Publishing Company, Inc.

Average Atomic Mass Avg. Atomic Mass weighted average of all isotopes on the Periodic Table round to 2 decimal places Avg. Atomic Mass

Average Atomic Mass EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O. Avg. Atomic Mass 16.00 amu

Average Atomic Mass EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. Avg. Atomic Mass 35.40 amu

Molar Mass of Compounds Use the periodic table to find the mass of the elements you need. Add up all the elements’ masses together appropriately. Example: H2 gas Mass of Hydrogen: 1.01 g/mole Subscript says I have 2 hydrogen atoms Mass of molecule = 1.01 + 1.01 = 2.02 g/mole

Molar Mass of Compounds Find the mass of ammonia, NH3 Find the mass of potassium sulfate, K2SO4

Ch. 10 - The Periodic Table II. Organization (p.286-291) Metallic Character Rows & Columns Table Sections

A. Metallic Character Metals Nonmetals Metalloids

B. Table Sections Representative Elements Transition Metals Inner Transition Metals

B. Table Sections Overall Configuration Lanthanides - part of period 6 Actinides - part of period 7

C. Columns & Rows Group (Family) Period

III. Periodic Trends (p.288-291) Terms Periodic Trends Dot Diagrams Ch. 10 - The Periodic Table III. Periodic Trends (p.288-291) Terms Periodic Trends Dot Diagrams

A. Terms Periodic Law Properties of elements repeat periodically when the elements are arranged by increasing atomic number.

A. Terms First Ionization Energy Valence Electrons e- in the outermost energy level Atomic Radius First Ionization Energy energy required to remove an e- from a neutral atom

B. Periodic Trends Atomic Radius Increases to the LEFT and DOWN.

B. Periodic Trends Increases to the RIGHT and UP. First Ionization Energy Increases to the RIGHT and UP.

Be or Ba Ca or Br Ba Ca B. Periodic Trends Which atom has the larger radius? Be or Ba Ca or Br Ba Ca

N or Bi Ba or Ne N Ne B. Periodic Trends Which atom has the higher 1st I.E.? N or Bi Ba or Ne N Ne

B. Periodic Trends Group # = # of valence e- (except He) Families have similar reactivity. Period # = # of energy levels 1A 2A 3A 4A 5A 6A 7A 8A

C. Dot Diagrams Dots represent the valence e-. EX: Sodium EX: Chlorine