Chapters 4.3 & 5.1-3 Notes A. Atomic Review B. Ch. 4.3: Modern Atomic Theory C. Ch. 5.1: Organizing the Elements D. Ch. 5.2: The Modern Periodic Table E. Ch 5.3: Representative Groups

A. Models of the Atom, pg 114-15 1803 – Dalton Model: pictures atoms as tiny, indestructible particles, with no internal structure 1897 – Thomson: discovers e-, leading to his “plum-pudding” model. He pictures e- embedded in a sphere of positive electric charge 1911 – Rutherford: states that atom has a dense, positively charged nucleus. E- move randomly in the space around the nucleus 1913 – Bohr: e- move in spherical orbits at fixed distances from the nucleus 1926 – Electron Cloud: more dense where electron is likely to be

A. Atomic Review Subatomic particles (p, n, e) Symbols Atomic # Mass # Isotopes

Diagram of an atom p – Proton e – Electron n - Neutron Proton Quark Nucleus What do you now know about the location of electrons? Neutron p – Proton e – Electron n - Neutron Electron

B. Ch. 4.3: Modern Atomic Theory Energy levels: the possible energies that electrons in an atom can have (staircase analogy) Electron Cloud: visual model of the most likely locations for electrons in an atom (windmill analogy) Orbital: a region of space around the nucleus where an electron is likely to be found (map of dots analogy) The most stable electron configuration is when electrons are in orbitals with the lowest possible energies.

Energy Levels Outer level electrons have more energy. Different levels hold different numbers of electrons.

Energy Level 3 18 electrons Energy Level 2 8 electrons Energy Level 1 2 electrons Nucleus

C. Ch 5.1: Organizing the Elements How did Mendeleev use solitaire to help him organize 63 elements? Compare and contrast Mendeleev’s periodic table to the table we use today. Mendeleev arranged elements by increasing atomic mass Today elements are arranged by increasing atomic number and changes in physical and chemical properties.

D. Ch. 5.2: The Modern Periodic Table What information can we get from the table without prior knowledge? Solids, liquids, gases, not found in nature, metals, nonmetals, metalloids… **NOTICE THE BOTTOM Lanthanide series and Actinide series Lanthanide can be found naturally on Earth, while the Actinide series only has some found in nature. The rest are made in labs only. All Actinide are radioactive. Characteristics of metals, nonmetals, and metalloids? Metals: hard, dense, high melting point, shiny, malleable, ductile, reactive, 88 elements to left of stairstep are metals or metal like, alloys (mixing of metals), corrosion Nonmetals: dull, brittle, not ductile or malleable, low density, low melting point, poor conductors, form compounds with each other Metalloids: characteristics of both, solids, shiny or dull, ductile, malleable, conductors (not well but can), semiconductors (conducts at high temps but not at low temps)

Families (Groups) and Periods (Rows) Ended here

Group or Family Trends Up and down; columns Elements in the same group have similar properties (have same # of valence electrons). Group 1, very reactive (only 1 valence electron) Group 8A, not very reactive (has all 8 valence electrons) Alkali metal, alkaline earth metal, halogens, noble gases… Start here 1/8

Row or Period Trends As you go to the right, you add 1 proton and 1 electron. The row ends when there are 8 electrons in the outer energy level Increasing Energy Levels

Variation Across a Period From left to right, the elements become less metallic and more nonmetallic in their properties. Na – reacts quickly and violently with water Mg – reacts with hot water Al – doesn’t react with water, but reacts w/oxygen Si – least reactive (except Ar) P & S – normally don’t react w/water, but react w/O and Cl, which is a highly reactive nonmetal. Ar – hardly reacts at all

SOME TRENDS OF THE PERIODIC TABLE We won’t work with all of these…

Periodic Table Trends Electronegativity – measure of attraction for electrons Ex: High electronegativity = strong attraction Ionization energy – energy required to remove an electron Ex: Low ionization E = easy to remove electron Electron affinity – ability of an atom to accept an electron Ex: High electron affinity = good at accepting e- Metallic character – same as atomic radii trend *Atomic radius – see next slide * Reactivity *Oxidation numbers – shows # of e- gained or lost *Valence electrons – see group number *Energy levels – see row number (* = characteristics you need to know)

Atomic Radii Pattern

Reason for atomic radii pattern? As you move down a group, the radii increases… The number of energy levels increases, and each subsequent energy level is further from the nucleus than the last As you move across a row, the radii decreases… More electrons are added to the same energy level, and more protons are added to the nucleus. Thus, there is a stronger force of attraction pulling the electrons closer to the nucleus. End here 1/6/16 Wednesday

How to draw an Electron Dot Diagram Find element on table. Write symbol. Look at the group number to determine the number of dots. Draw dots. Book discusses this in Chapter 6.1 Start here 1/7/16 Thursday

Electron Dot Diagrams Aka lewis dot structure Write steps explaining how to draw a Lewis Dot Structure.

E. Ch. 5.3: Representative Groups Valence Electron – electron that is in the highest occupied energy level of an atom (outermost energy level) Elements in a group have similar properties because they have the same number of valence e- The number of valence electrons help us predict how atoms will interact with each other (bonding).

Representative Groups, cont. Group 1A – Alkali Metals 1 valence electron, very reactive Group 2A – Alkaline Earth Metals 2 valence electrons, harder than alkali Group 4A – Carbon Family 4 valence electrons, metal, nonmetal, metalloid Group 5A – Nitrogen Family 5 valence electrons Group 6A – Oxygen Family 6 valence electrons Group 7A – Halogens 7 valence electrons, highly reactive nonmetals Group 8A – Noble Gases 8 valence electrons, colorless, odorless, unreactive