2 Chapter Outline History of chemistry (2.1 – 2.4) Chemical laws – start with slide 8 for content, 3-7 FYI slidesPath to the atomModern atomic structure (2.5)Molecules vs. Ions (2.6)Naming molecular and ionic compounds (2.8)Introduction to the periodic table (2.7)
3 History of ChemistryGreek Philosophers: 5th Century BCE (BCE = before the common era - replaces BC)The Greek philosophers were the first to reflect on the nature of matter. They proposed that all matter is made out of first 4 elements -- earth, air , water, fire. Aristotle added a fifth element, plasma (also called ether).
4 Greek PhilosophersDemocritus had an alternate view of matter. He proposed that matter was made up of tiny particles called atoms. His "theory" was not well accepted at the time.
5 Alchemy Alchemy: ~600-1600's CE (CE = common era, replaces AD)Alchemy developed at about the same time in China, India, and Greece. It spread into Europe in the 8th century.
6 Alchemy Alchemists had two pursuits Search for a means to convert “base” metals into goldSearch for the elixir of lifeSubstance that would lead to immortality
7 Alchemy Advances from Alchemy Many new substances where identified Plaster of Paris, nitric acid….New lab techniques and equipment developedNew medicines identified
8 Modern Chemistry, ~1600 onFirst chemists/physicists to use scientific methodBoyle - elementsLavoisier – law of conservation of matterProust - law of definite proportionsDalton – law of multiple proportions, atomic theoryAvogadro - hypothesisThomson – charge to mass ratio for an electronMillikan – charge on the electronBequerel and the Curies - radioactivityRutherford – nuclear atom
9 Modern Chemistry Robert Boyle: ~1660 Proposed a substance to be an element unless it can be broken down into simpler substances.Proposed one of the gas laws – CH 5Boyle – studied the relationship between pressure and volume of a gasBoyle still believed in alchemy – that metals could be converted into gold
10 Lavoisier: ~1760 Law of Conservation of Matter Matter is neither created nor destroyed in a chemical reaction.Called the father of modern chemistryBeheaded for his political believes
11 Proust: late 1700s Law of Definite Composition/Proportions A given compound always contains the same proportion of elements by mass.
12 Law of Multiple Proportions John Dalton: ~1800Law of Multiple ProportionsWhen two elements form more than one compound, the ratios of the masses of the second element that combines with one gram of the first element can always be reduced to small whole numbers.
13 Law of Multiple Proportions Example Consider two g samples 2 different compounds containing only C and H. Compound A: 27.2 g of CCompound B: 42.9 g of C.Show how this data illustrates the law of multiple proportions.
14 Dalton also proposed the first table of atomic masses Most masses later need revisionDalton is best known for proposing Atomic Theory
15 Dalton’s Atomic Theory Elements are made up of tiny particles called atoms.Atoms are indivisible and indestructibleAtoms of a given element are identicalAtoms of different elements differ in some fundamental way(s)
16 Dalton’s Atomic Theory Compounds form when atoms of different elements combine with each other.A given compound always has the same relative number and types of elements.
17 Dalton’s Atomic Theory Chemical reactions occur when atoms change how they are bound to each other.Individual atoms are not changed, just rearranged
18 Avogadro: 1811 Avogadro's Hypothesis At the same temperature and pressure equal volumes of gases contain the same number of particles.Based on Guy-Lussac’s dataSee page 41/42
19 From Dalton to Atomic Structure Dalton’s atomic theory lead to much research on the nature of the atom.This research showed the atom to made up of smaller particles.
20 J.J. ThomsonThomson measured the deflection of a cathode ray beam in electrical and magnetic fields of known strengths.Applied electrical fieldCathode ray+(+)Metal electrode(-)Metal electrode-Cathode ray tube experiment, pg 43
21 Thomson found the cathode rays were attracted by the positive charge and repelled by negative These findings clearly indicated that the rays consisted of negatively charged particles.Today we know these particles as electrons.
22 Thomson measured the deflection of the beam in a magnetic field and more! From his data he determined the charge:mass ratio for an electrone = x 108 C/gme = charge on the electron in coulombsM = mass of an electron in gramsCathode ray tube experiment
23 Thomson also found that the cathode ray particles were identical regardless of source. Concluded all elements contain these negative particles (electrons)
24 J.J. ThomsonThomson:identified cathode ray beams as a stream of negatively charged particlescalculated the charge to mass ratio for these negatively charged particlesproposed the existence of positively charged particlesTo balance the negative charge of the electrons
25 Millikan ~1909Millikan’s oil drop experiment allowed him to determine the charge on an electronThis charge can be plugged into Thomson’s formula and the mass of the electron calculatedMass electron = x kgPage 44
26 Radioactivity Becquerel, Marie and Pierre Curie: ~1896 Henri Becquerel - observed the natural emission of energy/rays by uranium.Marie and Pierre Curie studied “Becquerel's rays”.The Curies’ findings suggested that matter was composed of smaller particles than atoms.The Curies coined the term radioactivity to describe the rays emitted.
27 Radioactivity Three types of radioactivity were identified: gamma rays - very high energy lightbeta particles - high energy electronsalpha particles - He+2 particles2 protons and 2 neutrons
28 Plum Pudding Model of the Atom In the early 1900’s the accepted model of the atom was called the plum pudding model of the atomElectrons (tiny and negatively charged) were pictured to be dispersed in a ‘cloud’ of positive charge.Proposed by JJ Thomson and Lord Kelvin in 1904
29 Rutherford and the Nuclear Atom In 1911 an experiment conducted in Ernest Rutherford’s lab showed the “plum pudding” model to be incorrect.Experiment was conducted by Geiger and Marsden and the findings interpreted by Rutherford.See page 45
30 Rutherford’s Atom First to propose a nuclear atom. An atom has a dense positive center containing all of positive charge and most of the mass of the atom – the nucleusElectrons are scattered in the empty space around the nucleusElectrons occupy a volume that is huge as compared to the size of the nucleus.
31 A New Model of the Atom Expected based on Plum pudding model Rutherford’s modelBased on ”his” results
32 Modern Atomic Structure Rutherford continued to study the atom and the positive matter of the atom.1919, + particle named the proton~1932 James Chadwick proposed the existence of a third subatomic particle, the neutron.
33 Subatomic Particles Subatomic Particle Charge Mass, amu Location in atomElectron(e-)-10 amuOutside of nucleusProton (p)+1~1 amuNucleusNeutron (n)
34 Mass of Subatomic Particles Protons and neutrons have ~ the same mass (in the range of kg).Mass of each and of individual atoms is often expressed in amu rather than gramsAtomic mass unit (amu) – 1/12 the mass of a carbon-12 atom
35 Mass of Subatomic Particles The mass of the electron (10-31 kg) is tiny as compared to that of the proton and neutron (10-27 kg) .Therefore, the electron’s mass is considered to be ~0 amu when calculating the mass of an atom.
36 Subatomic Particles and the Elements Each element has a unique number of protons.Number of protons defines the element.Atomic # = # protons6C
37 Subatomic ParticlesSince atoms are neutral, for every proton there is a/n _________.When atoms interact to form compounds, it is their ___________ that interact.
38 TermsMass number = sum of the # of protons and the # neutrons in the nucleus of an atomFOR MOST ELEMENTS THE MASS NUMBER IF NOT ON THE PERIODIC TABLE.You will be given enough information to determine mass number or number of neutrons.
39 Terms Isotopes = atoms of a given element that differ in mass number Isotopes have the same number of _____________.Isotopes differ in the number of _______.
40 B Isotopes Writing atomic symbols for isotopes pg 46 11 Mass # Symbol for element5Atomic #
41 FAQ - Isotopes C Po When is mass number found on the periodic table? What’s the atomic mass? Is it the same as the mass number?C(209)Po
42 Molecules and Ions (2.6)Atoms of different elements combine to form compoundsAtoms in compounds are held together by chemical bonds.Bonds involve interactions of the bonding atoms’ ________
43 Bonding There are two types of bonds: Covalent bonds – bonding atoms share electronsAtoms are always nonmetal atomsCovalently bonded atoms form moleculesWays to represent moleculesChemical formula; H2OStructural formulaOH H
44 Bonding Ionic bonds – attractive force among oppositely charged ions Bond formed between metal cations and nonmetal anionsNo molecules involved
45 Ions - Terms Ion – charged atom or group of atoms Formed when atoms gain or lose electronsCation – positively charged ionFormed when an atom _______ electronsAnion – negatively charged ionFormed when an atom ______ electrons
46 IonsDescribing ion formationCation example:Anion example:
48 Types of Binary Compounds Type I binary ionic compoundsMetal forms only one ionType II binary ionic compoundsMetal forms more than one ionUse roman numerals to indicate the charge on the ionType III binary covalent compoundsCompound between 2 nonmetals
49 Types I Binary Compounds Compound between a metal and a nonmetalMetal forms only one ionName the cation and then the anion.Name of the cation is the name of the elementName of the anion is the name of the nonmetal with the ending changed to “ide”
50 Monoatomic cations to know Group #Charge on ionexamplesIA+1Na1+ sodium (ion)K potassium (ion)IIA+2Mg2+ magnesium (ion)IIIA metals+3Al3+ aluminum (ion)
51 Monoatomic anions to know Group #Charge on ionexamplesVA-3N3- nitride (ion)P3- phosphide (ion)VIA-2O2- oxide (ion)S sulfideVIIA-1F1- fluoride (ion)Cl1- chloride (ion)Br1- bromide (ion)I iodide (ion)
52 PracticeName chemical formulaChemical formula name