Atoms, Molecules, and Ions

Chemistry Timeline #1 B.C. 400 B.C. Demokritos and Leucippos use the term "atomos” 1500's  Georg Bauer: systematic metallurgy  Paracelsus: medicinal application of minerals 1600's Robert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of elements 1700s'  Georg Stahl: Phlogiston Theory  Joseph Priestly: Discovery of oxygen  Antoine Lavoisier: The role of oxygen in combustion, law of conservation of mass, first modern chemistry textbook  2000 years of Alchemy

Chemistry Timeline #2 1800's  Joseph Proust: The law of definite proportion (composition)  John Dalton: The Atomic Theory, The law of multiple proportions  Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules  Amadeo Avogadro: Molar volumes of gases  Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements  Dmitri Mendeleyev: The periodic table  J.J. Thomson: discovery of the electron  Henri Becquerel: Discovery of radioactivity 1900's  Robert Millikan: Charge and mass of the electron  Ernest Rutherford: Existence of the nucleus, and its relative size  Meitner & Fermi: Sustained nuclear fission  Ernest Lawrence: The cyclotron and trans-uranium elements

Dalton’s Atomic Theory (1808)  Atoms cannot be subdivided, created, or destroyed  Atoms of different elements combine in simple whole-number ratios to form chemical compounds  In chemical reactions, atoms are combined, separated, or rearranged  All matter is composed of extremely small particles called atoms  Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties John Dalton

Modern Atomic Theory Several changes have been made to Dalton’s theory. Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element.

Modern Atomic Theory #2 Dalton said: Modern theory states: Atoms cannot be subdivided, created, or destroyed Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

Mass of the Electron 1909 – Robert Millikan determines the mass of the electron. The oil drop apparatus Mass of the electron is x kg

Conclusions from the Study of the Electron  Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.  Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons  Electrons have so little mass that atoms must contain other particles that account for most of the mass

Rutherford’s Gold Foil Experiment  Alpha particles are helium nuclei  Particles were fired at a thin sheet of gold foil  Particle hits on the detecting screen (film) are recorded

Try it Yourself! In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

The Answers Target #1Target #2

Rutherford’s Findings  The nucleus is small  The nucleus is dense  The nucleus is positively charged  Most of the particles passed right through  A few particles were deflected  VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions:

Atomic Particles ParticleChargeMass (kg)Location Electron9.109 x Electron cloud Proton x Nucleus Neutron x Nucleus

The Atomic Scale  Most of the mass of the atom is in the nucleus (protons and neutrons)  Electrons are found outside of the nucleus (the electron cloud)  Most of the volume of the atom is empty space “q” is a particle called a “quark”

About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”

Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) 110 Hydrogen-2 (deuterium) 111 Hydrogen-3 (tritium) 112

Atomic Masses IsotopeSymbolComposition of the nucleus % in nature Carbon C6 protons 6 neutrons 98.89% Carbon C6 protons 7 neutrons 1.11% Carbon C6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally isotopes of that element. Carbon =

Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element# of protonsAtomic # (Z) Carbon66 Phosphorus15 Gold79

Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p + + n 0 Nuclidep+p+ n0n0 e-e- Mass # Oxygen Arsenic Phosphorus153116

Molecules Two or more atoms of the same or different elements, covalently bonded together. Molecules are discrete structures, and their formulas represent each atom present in the molecule. Benzene, C 6 H 6

Covalent Network Substances Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Why Not?? GraphiteDiamond

IonsIons  Cation: A positive ion Mg 2+, NH 4 +Mg 2+, NH 4 +  Anion: A negative ion  Cl , SO 4 2   Ionic Bonding: Force of attraction between oppositely charged ions.  Ionic compounds form crystals, so their formulas are written empirically (lowest whole number ratio of ions).

Periodic Table with Group Names

Easily lose valence electron (Reducing agents) React violently with water Large hydration energy React with halogens to form salts The Properties of a Group: the Alkali Metals

Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+H+H+H+ Li + Na + K+K+K+K+

Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

Predicting Ionic Charges Group 13: Loses 3 Loses 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

Predicting Ionic Charges Group 14: Loses 4 Loses 4 electrons or gains 4 electrons Caution! C 2 2- and C 4- are both called carbide

Predicting Ionic Charges Group 15: Gains 3 Gains 3 electrons to form 3- ions N 3- P 3- As 3- Nitride Phosphide Arsenide

Predicting Ionic Charges Group 16: Gains 2 Gains 2 electrons to form 2- ions O 2- S 2- Se 2- Oxide Sulfide Selenide

Predicting Ionic Charges Group 17: Gains 1 Gains 1 electron to form 1- ions F 1- Cl 1- Br 1- Fluoride Chloride Bromide I 1- Iodide

Predicting Ionic Charges Group 18: Stable Noble gases do not form ions! Stable Noble gases do not form ions!

Predicting Ionic Charges Groups : Many transition elements Many transition elements have more than one possible oxidation state. have more than one possible oxidation state. Iron(II) = Fe 2+ Iron(III) = Fe 3+

Predicting Ionic Charges Groups : Some transition elements Some transition elements have only one possible oxidation state. have only one possible oxidation state. Zinc = Zn 2+ Silver = Ag +

Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

Writing Ionic Compound Formulas Example: Iron(III) chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe 3+ Cl - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 3

Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 23

Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO Check to see if charges are balanced. They are balanced!

Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! Zn 2+ OH - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ PO Check to see if charges are balanced. They ARE balanced!

Naming Ionic Compounds 1. Cation first, then anion1. Cation first, then anion 2. Monatomic cation = name of the element2. Monatomic cation = name of the element Ca 2+ = calcium ionCa 2+ = calcium ion 3. Monatomic anion = root + -ide3. Monatomic anion = root + -ide Cl  = chlorideCl  = chloride CaCl 2 = calcium chlorideCaCl 2 = calcium chloride

Naming Ionic Compounds (continued)  some metal forms more than one cation  use Roman numeral in name PbCl 2 Pb 2+ is the lead(II) cation PbCl 2 = lead(II) chloride Metals with multiple oxidation states

Naming Binary Compounds  Compounds between two nonmetals  First element in the formula is named first.  Second element is named as if it were an anion.  Use prefixes  Only use mono on second element - P 2 O 5 = CO 2 = CO = N 2 O = diphosphorus pentoxide carbon dioxide carbon monoxide dinitrogen monoxide