Atoms and the Periodic Table

Atoms  Atoms are the smallest pieces of matter that contain all the properties of a specific element  Each element contains only one type of atom

Atomic Models  Have been revised many times to explain new discoveries  Democritus (4 th century B.C.) thought all matter was made of particles he called the atom  Theory was modified when subatomic particles were discovered

Atomic Models  Plum Pudding Model  (1904) developed by J.J. Thomson  Planetary Model  (1911) developed by Ernest Rutherford

Newer Models  Bohr’s Model (1913)  Developed by Niels Bohr  Electron Cloud Model (1925) 

Inside the Atom  Atoms are made up of smaller particles  These particles are found in different regions of the atom

Protons (p + )  Positively charged particles found in nucleus of atom  Have an electrical charge of +1  Mass of 1 a.m.u.  Composed of quarks  Discovered by Ernest Rutherford using Gold Foil Experiment

Protons  The number of protons found in the nucleus of an atom determines the identity of the atom  All oxygen atoms have 8 protons, all uranium atoms have 92 protons  If the number of protons change the identity of the atom changes.

Neutrons (n 0 )  Neutral particles found in nucleus of atom  Have no electrical charge  Mass of 1 a.m.u.  Composed of quarks  Discovered by James Chadwick

Nucleus  The nucleus is the positively charged dense core in the center of the atom  Houses protons and neutrons  Contains 99.9% of mass of atom  Extremely small compare to the entire size of the atom

Electrons (e - )  Negatively charged particles found in electron cloud  Have an electrical charge of -1  Constantly moving around outside nucleus  Have essentially no mass  Discovered by J.J. Thomson during cathode ray experiment

Electrons  The number of electrons in a neutral atom is equal to the number of protons in the atom  Neutral oxygen has 8 protons, therefore it has 8 electrons  Neutral lead has 82 protons, therefore, it has 82 electrons

Electron Cloud  Region around nucleus  Houses electrons

Valence Electrons  Electrons in the outermost energy level of an electron are called valence electrons  These are the electrons furthest from the nucleus

Symbols  Elements are listed by their chemical symbols  Symbols are usually either one capital letter like C for Carbon, or one capital and one lowercase letter like Ne for Neon

Periodic Table  The periodic table gives much information we need to learn more about the atom of each element

Atomic Number (Z) Whole number shown on periodic table  Periodic table is arranged by atomic number Atomic Number = # of Protons *Also gives the number of electrons if the atom is neutral

Atomic Number

Mass Number (A)  The mass number is the sum of the total number of protons and neutrons in the atom  Mass # = # p + + # n 0

Atom Math Atomic Number Symbol Element Name Atomic Mass

Practice Zn Oxygen Lithium # n 0 = Mass # - Atomic #

Isotopes  Isotopes are atoms of the same element that have different numbers of neutrons  All atoms are isotopes  Each element has isotopes that are more common than others

Nuclear Symbol  Isotopes can be designated with their nuclear symbol

Hyphen Notation  Isotopes can also be designated using hyphen notation Carbon-16 Element NameMass Number

Write the Nuclear Symbol and Hyphen Notation for the Following Isotopes  Lithium isotope with 3 protons and 4 neutrons  Sulfur isotope with 17 neutrons  Lead with 122 neutrons

Ions  Ions are atoms or groups of atoms that have a net positive or negative charge  The charge results from an unequal number of electrons and protons within atom of group of atoms

Ions  Ions with more electrons than protons have a negative charge  For each extra electron the negative charge increases by one  Ions with less electrons than protons have a positive charge  For each missing electron the positive charge increases by one

Ions  The charge of an ion is indicated on its element symbol with a positive or negative sign on the upper right side and number equal to the magnitude of the charge  The number one is not included

Ions  F -  9 protons – 10 electrons = -1 charge  Ca 2+  20 protons – 18 electrons = -2 charge  P 3-  15 protons – 18 electrons = -3 charge

Common Ions (Need to memorize these)  Many of the main group elements in the same column form ions with same charge  Column 1 (Li, Na, K, Rb, Cs): +1  Column 2 (Be, Mg, Ca, Sr, Ba): +2  Column 13 (Al, Ga): +3  Column 15 (N, P, As): -3  Column 16 (O, S, Se, Te): -2  Column 17 (F, Cl, Br, I): -1

Atomic Mass  The average atomic mass is the number at the bottom of this square  Found by averaging the natural abundances of its isotopes

Calculating Average Atomic Mass (amu) If abundance is given as percent value:  If abundance is given as decimal value:

Average Atomic Mass Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?

Uranium has three common isotopes. If the abundance of 234U is , the abundance of 235U is , and the abundance of 238U is.9928, what is the average atomic mass of uranium?

Electron Configuration  An electron configuration is a shorthand description of how electrons are arranged around the nucleus of an atom.  Electrons generally do not orbit the nucleus rather, they move in certain patterns based on three factors: (1) their energies, (2) their orientations in space and (3) their angular momentum

A. General Rules  Pauli Exclusion Principle  Each orbital can hold TWO electrons with opposite spins.

A. General Rules  Aufbau Principle  Electrons fill the lowest energy orbitals first.  “Lazy Tenant Rule”

O 8e -  Orbital Diagram  Electron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

RIGHT WRONG A. General Rules  Hund’s Rule  Within a sublevel, place one e - per orbital before pairing them.  “Empty Bus Seat Rule”

 Shorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 B. Notation  Longhand Configuration

© 1998 by Harcourt Brace & Company s p d (n-1) f (n-2) C. Periodic Patterns

 Period #  energy level (subtract for d & f)  A/B Group #  total # of valence e -  Column within sublevel block  # of e - in sublevel

s-block1st Period 1s 1 1st column of s-block C. Periodic Patterns  Example - Hydrogen

C. Periodic Patterns  Shorthand Configuration  Core e - : Go up one row and over to the Noble Gas.  Valence e - : On the next row, fill in the # of e - in each sublevel.

[Ar]4s 2 3d 10 4p 2 C. Periodic Patterns  Example - Germanium

 Full energy level  Full sublevel (s, p, d, f)  Half-full sublevel D. Stability

 Electron Configuration Exceptions  Copper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10  Copper gains stability with a full d-sublevel. D. Stability

 Electron Configuration Exceptions  Chromium EXPECT :[Ar] 4s 2 3d 4 ACTUALLY :[Ar] 4s 1 3d 5  Chromium gains stability with a half-full d- sublevel. D. Stability

 Ion Formation  Atoms gain or lose electrons to become more stable.  Isoelectronic with the Noble Gases.

O 2- 10e - [He] 2s 2 2p 6 D. Stability  Ion Electron Configuration  Write the e - config for the closest Noble Gas  EX: Oxygen ion  O 2-  Ne

D. Lewis Diagrams  Also called electron dot diagrams  Dots represent the valence e -  Ex: Sodium  Ex: Chlorine Lewis Diagram for Oxygen

Steps to Draw Lewis Structures 1.Determine how many valence are in the element. 2.Staring on the Right side of the element draw a dot to represent a valence electron. 3.Place one dot on each side of the symbol. One electron must be drawn around each side of the element before a second electron can be added to any side.

Steps to Draw Lewis Structures