Atoms, Molecules and Ions

Dalton’s Atomic Theory (1808) 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 Atoms cannot be subdivided, created, or destroyed (Law of conservation of mass) Atoms of different elements combine in simple whole-number ratios to form chemical compounds (Law of constant composition) In chemical reactions, atoms are combined, separated, or rearranged John Dalton

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 is 9.109 x 10-31 kg 1909 – Robert Millikan determines the mass of the electron. Mass of the electron is 9.109 x 10-31 kg The oil drop apparatus

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

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

Atomic Particles Particle Charge Mass (kg) Location Electron -1 9.109 x 10-31 (1/1836 amu) Electron cloud Proton +1 1.673 x 10-27 (1 amu) Nucleus Neutron 1.675 x 10-27 (1 amu)

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

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

Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon 6 Phosphorous 15 Gold 79

Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 Nuclide p+ n0 e- Mass # Oxygen – 18 8 10 18 Arsenic – 75 33 75 Phosphorus – 31 15 16

Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Nucleus Hydrogen-1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2

Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011 Carbon = 12.011 Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 7 neutrons 1.11% Carbon-14 14C 8 neutrons <0.01%

Calculating Average Atomic Mass Average atomic mass = S(% of each isotope)(mass of each isotope) 100 Example: If 69Ga and 71Ga occur in the percents 62.1 and 37.9, calculate the average atomic mass of gallium atoms.

Molecules Two or more atoms of the same or different elements, ___________ bonded together. Molecules are discrete structures, and their formulas represent each atom present in the molecule. Benzene, C6H6

Examples of Molecules Examples of Molecules – Substance and Formula Element or Compound Description Hydrogen (H2); Oxygen (O2); Nitrogen (N2); Fluorine (F2); Chlorine (Cl2); Iodine (I2); Bromine (Br2) Elements Diatomic Water (H2O) Compound Polyatomic Ammonia (NH3)

Covalent Network Substances Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Groups of these substances are called __________. Why Not?? Graphite Diamond

Ions Atoms have equal numbers of protons and electrons so they have ________ When atoms gain or lose electrons, the proton:electron numbers are uneven causing a charged particle called an ____

Ions Cation: A ________ ion Anion: A _______ ion Mg2+ (monatomic), NH4+ (polyatomic) Anion: A _______ ion Cl- (monatomic), SO42- (polyatomic) Ionic Bonding: Force of attraction between oppositely charged ions. Ionic compounds form crystals, so their formulas are written _________ (lowest whole number ratio of ions).

Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+ Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Sr2+ Ba2+

Predicting Ionic Charges Group 13: Loses 3 electrons to form 3+ ions B3+ Al3+ Ga3+ Group 14: Loses 4 electrons or gains 4 electrons Caution! C22- and C4- are both called carbide

Predicting Ionic Charges Group 15: Gains 3 electrons to form 3- ions N3- P3- As3- Nitride Phosphide Arsenide Group 16: Gains 2 electrons to form 2- ions O2- S2- Se2- Oxide Sulfide Selenide

Predicting Ionic Charges Group 17: Gains 1 electron to form 1- ions F1- Cl1- Br1- I1- Fluoride Chloride Bromide Iodide Group 18: Stable Noble gases do not form ions!

Predicting Ionic Charges Groups 3-12: Many transition elements have more than one possible oxidation state (charge). Iron(II) = Fe2+ Iron(III) = Fe3+ Some transition elements have only one possible oxidation state. Zinc = Zn2+ Silver = Ag+

Binary Compounds of Nonmetals and Metals (Ionic Compounds) To write the formula from the name Write the formulas for the _____ and _____, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges , if necessary, using ________. Use parentheses if you need more than one of a polyatomic ion.

Writing Ionic Compound Formulas Example: Barium nitrate Example: Ammonium sulfate Example: Iron(III) chloride Example: Aluminum sulfide Example: Magnesium carbonate Example: Zinc hydroxide Example: Aluminum phosphate

Naming Ionic Compounds 1. Cation first, then anion 2. Monatomic cation = name of the element Ca2+ = calcium ion 3. Monatomic anion = root + -ide Cl- = chloride CaCl2 = calcium chloride

Naming Ionic Compounds Metals with multiple oxidation states Some metal forms more than one cation Use Roman numeral in name PbCl2 Pb2+ is the lead(II) cation PbCl2 = lead(II) chloride

Binary Acids An acid is a compound that produces hydrogen ions in water Formed when hydrogen ions combine with monatomic anions To name an acid Use the prefix hydro followed by the non-metal name modified to an –ic ending Add the word acid Example: HCl and HI

Polyatomic Ions and Oxyanions Polyatomic ion – more than one element are combined to create a species with a charge Have to memorize these Oxyanions – a polyatomic ion that contains oxygen To name a compound containing a polyatomic ion Positive ion written first Polyatomic ion written second, unmodified Example: K2CO3

Oxyacids Formed when hydrogen ions combine with polyatomic oxyanions To name an oxyacid Use the name of the oxyanion and replace the –ite ending with –ous or the –ate ending with –ic Add the word acid Example: HClO4 and HNO3

Naming Binary Compounds of Two Nonmetals Two nonmetals form molecules First element in the formula is named first. Second element is named as if it were an anion. Use prefixes – mono, di, tri, tetra, penta, hexa, hepta and octa Only use mono on second element - P2O5 = diphosphorus pentoxide CO2 = carbon dioxide CO =carbon monoxide N2O = dinitrogen monoxide

Hydrates Ionic formula units with water associated with them Water molecules are incorporated into the solid structure of the ions To name a hydrate Use the naming of an ionic compound followed by the term hydrate with an appropriate prefix for the hydrate Example: CuSO4.5H2O