Unit 2

 There are about 118 different known elements with 88 of them occurring naturally.  The names and symbols of each element are listed on the periodic table in your book.

 Names for elements come from many different sources. Many have Latin or Greek origins However, more recent discoveries are named for descriptions, Famous Scientists, or Place of discovery  We use abbreviations to simplify when writing called Element Symbols.

 Examples: Berkelium (Bk)- named for place of discovery- Berkeley, California Copper (Cu)- Latin- cuprum or cyprium, discovered in Cyprus Lead (Pb)- Latin,- plumbum, meaning heavy Oxygen (O)- French, oxygene, generator of acid, derived from the Greek, oxy and gene meaning acid forming  Oxygen was thought to be part of all acids, but it’s not. Can you guess which element is part of all acids?

 Element Symbols- First or first two letters of the element names. The first letter is always capitalized The second letter is always lowercase  O- oxygen, C- carbon, Ne- neon, Si- silicon

 Element Symbols- Sometimes the first two letters are not the first two letters of the name. Symbols are from the old names Symbols are from other letters in the name  Gold- Aurum- Au  Lead- Plumbum- Pb  Zinc- Zn  Cadmium- Cd

 Law of Constant Composition: A given compound always contains elements in exactly the same proportion by mass.  This observation along with others became the basis for Dalton’s Atomic Theory.

 1. Elements are made of tiny particles called Atoms. All matter is made of atoms. Atoms are indivisible and indestructible particles.  2. All atoms of a given element are identical, both in mass and in properties.  3. The atoms of a given element are different from those of any other element.

 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative number and types of atoms.  5. Atoms are the units of chemical change. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.

 Chemical Formula: a way of writing a compound using only symbols and numbers. The atoms are indicated by he element symbols and the number of each type of atom is indicated by a subscript.  Water- H 2 O Water contains 2 hydrogen atoms and 1 oxygen atom  Carbon dioxide- CO 2 Carbon dioxide contains 1 carbon atom and 2 oxygen atoms

 The atom is composed of 3 subatomic particles called:  1. Protons (p + )  2. Neutrons (n 0 )  3. Electrons (e - )  Protons and Electrons are always equal in number in neutral atoms

 Nuclear Atom: An atom with a dense center of positive charged around which tiny electrons move in a space that was otherwise empty.  Nucleus: The relatively small, dense center of positive charge in an atom. Made of Protons and Neutrons

 Proton: A positively charged subatomic particle located in the atomic nucleus.  Neutron: A subatomic particle with no charge located in the atomic nucleus.  Electron: A negatively charged subatomic particle located outside of the nucleus.

ParticleSymbolRelative Charge Relative Mass Electrone-1-1 amu Protonp amu NeutronN01839 amu

 Atomic Number: The number of protons intside the nucleus of an atom. Generally given the symbol Z  Mass Number: The sum of the number of protons and neutrons in the nucleus of an atom. Generally given the symbol A  The element symbol is given the symbol X A Z X = Na 23 11

 In natural samples of an element you may find atoms with different masses. This phenomena can be explained by isotopes  Isotopes: Atoms with the same number of protons, but with different numbers of neutrons.’  Most atoms have at least two stable isotopes  Exceptions:  Aluminum, Fluorine, and Phosphorus have only 1  Tin has 10

 When we refer to an isotope we use its name and mass number We don’t have to give the atomic number because it is the same in all isotopes of a given element  Example: Boron-10 ( 10 B) Boron- 11 ( 11 B)

 Hydrogen is an exception to the name and mass number rule The isotopes of hydrogen are so important that they have special symbols and names  Protium (P)- hydrogen with no neutrons  Deuterium (D)- hydrogen with 1 neutron  Tritium (T)- hydrogen with 2 neutrons

 Percent abundance of each isotope can be calculated if the masses of the isotopes is found using a mass spectrometer  Using a mass spectrometer we find that the mass of 10 B is amu and 11 B is amu

 10 B is amu  11 B is amu  Average Atomic Mass is amu  By looking at the information, which isotope occurs in the greatest abundance?

 = (% 10 B ) + (% 11 B )  We know from algebra that when you add two percents they must equal 100 We reduce this to two decimals equal to 1  % 10 B + % 11 B = 1, where % 10 B = x and % 11 B = y  So, x + y = 1  We need to have the equation in terms of one variable so, y = 1 - x  = (x ) + ((1 – x) )

 We find that 10 B has an abundance of 19.91% and 11 B has an abundance of 80.09%  This means that in an average natural sample of 10,000 boron atoms you would find that 1,991 would be 10 B atoms and 8,009 would be 11 B atoms

 Antimony, Sb, has two stable isotopes with experimentally determined masses of amu ( 121 Sb) and amu ( 123 Sb). What are the relative abundances of these isotopes ?

 Periodic Table: A chart that shows all the known elements and gives you information about each one. Elements are listed on the periodic table in order of increasing atomic number One of the most useful tools in chemistry

 Elements are arranged in vertical Groups and horizontal Periods.  Periodic tables used in the United States have groups numbered 1-8 followed by the letter: A or B.

 A groups are main group elements  B groups are transition elements  Group 1A- Alkali Metals  Group 2A- Alkaline Earth Metals  Groups 3B-12B – Transition Metals  Group 7A- Halogens  Group 8A- Noble Gases

 The horizontal periods are numbered from 1-7  Period 1 contains only H and He  Periods 2 & 3 contain 8 elements  Periods 4 & 5 contain 18 elements  Periods 6 & 7 contain 32 elements

 The table is split into 3 basic parts:  1. Metals  2. Non-metals  3. Metalloids/semimetals

 1. Metals:  high electrical conductivity  high luster (shininess)  high ductility (can be drawn into wires)  high malleability (can be rolled into sheets)  can form alloys (solutions of one or more metals in another metal)  All metals are solids except for Mercury

 2. Non-metals:  Nonlustrous  poor conductors of electricity  All lie to the right of the diagonal line that stretches from B to Te in the periodic table  Some are solids; bromine is a liquid, and a few, like nitrogen and oxygen are gases at room temp.

 3. Metalloids/ semimetals:  display characteristics of both metals and nonmetals.  Only silicon, germanium, arsenic, antimony, and tellurium are in this category

Natural States of the Elements  Most elements are reactive and are not found naturally in pure form  However, there are a few exceptions

 Gold, Silver, and Platinum are called Noble Metals because they are relatively unreactive  Group 8 elements are called Noble Gases because they do not combine readily with other elements. He, Ne, Ar, Kr, Xe, Rn

 Diatomic Molecules: Molecules made up of two atoms. H 2, O 2, N 2, Cl 2, F 2, Br 2, and I 2 are diatomic molecules in pure, elemental form.  All the elements of Group 7 are diatomic molecules

 Allotropes: Different forms of elements where there are the same atoms, but they are structured differently. Carbon comes as Diamond, Graphite, and Buckminsterfullerene.

 Ion: An atom or group of atoms that has a positive or negative charge.  Taking a neutral atom and adding or subtracting one or more electrons can result in a charged ion.

 Positive ions are called Cations. Produced when an electron is lost from a neutral atom.  Mg 2+ and Na + are examples of cations.  Magnesium normally has 12 protons and 12 electrons, but when 2 electrons are lost it becomes a cation with a 2+ charge.

Ions  Negative ions are called Anions. Produced when an electron is gained to a neutral atom.  Cl - and O 2- are examples of anions  Chlorine normally has 17 protons and 17 electrons, but when it gains an electron it becomes an anion with a 1- charge.

Ions  Individual atoms that have lost or gained electrons are called Monatomic ions  Na + 11 protons, 10 electrons  O 2- 8 protons, 10 electrons  Al protons, 10electrons

Ions  Groups 1A- 3A form positive ions with a charge equal to the group number of the metal  1A Na 1 electron lost Na +  2A Ca 2 electrons lost Ca 2+  3A Al 3 electrons lost Al 3+

Ions  Nonmetals often form ion with a negative charge equal to 8-(group #) of the element.  5A N 8-5= 3 electrons gained N 3-  6A S 8-6= 2 electrons gained S 2-  7A B 8-7= 1 electron gained B -

Ions  There is no easily predictable pattern for determining charges on transition metals  Many of them also form several different ions Iron can be Fe 2+ or Fe 3+

Ions  Polyatomic ions contain two or more atoms with the resulting compound having an electric charge.  NH 4 + Four hydrogen atoms surround a nitrogen atom, and the group has a 1+ charge.

Compounds That Contain Ions  Many chemical compounds contain ions. We know this because electric currents can run through them. Substances can only conduct electric current if the ions can move freely  Salt water has ions, however pure salt and pure water cannot conduct electricity.

 Ionic Compounds: A compound that results when a metal reacts with a nonmetal to form cations and anions. The result must have a net charge of zero.  1. Both positive and negative ions must be present.  2. The numbers of cations and anions must be such that the net charge is zero.

Compounds That Contain Ions  Na + + Cl - = NaCl  Mg 2+ + Cl - = MgCl 2  Li + + N 3- = Li 3 N