Presentation on theme: "Atoms and Elements Chapter 4"— Presentation transcript:
1 Atoms and Elements Chapter 4 Chem 10 Spring 2008 Chp 4Atoms and Elements Chapter 4Tro, 2nd ed.
2 Each element has a number. ELEMENTS AND ATOMSAll known substances on Earth and probably the universe are formed by combinations of more than 100 elements.An element is a fundamental or elementary substance that cannot be broken down into simpler substances by chemical means.Each element has a number.Beginning with hydrogen as 1, the elements are numbered in order of increasing complexity.
3 ELEMENTS AND ATOMSMost substances can be decomposed into two or more simpler substances.Water can be decomposed into hydrogen and oxygen.Table salt can be decomposed into sodium and chlorine.An element cannot be decomposed into a simpler substance.An atom is the smallest particle of an element that can exist.An atom is the smallest unit of an element that can enter into a chemical reaction.
5 The diameter of an atom is 0.1 to 0.5 nm. This is 1 to 5 ten billionths of a meter.Even smaller particles than atoms exist. These are called subatomic particles.If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot.
6 Dalton’s Atomic Theory Elements are composed of minute indivisible particles called atoms.Atoms under special circumstances can be decomposed.Modern research has demonstrated that atoms are composed of subatomic particles.Atoms of the same element are alike in mass and size. Atoms of different elements have different masses and sizes. (Is this still correct?)Chemical compounds are formed by the union of two or atoms of different elements in whole number ratios.
7 In 1897 Sir Joseph Thompson demonstrated that cathode rays: WHAT’S IN AN ATOM?DISCOVERING SUBATOMIC PARTICLES!In 1897 Sir Joseph Thompson demonstrated that cathode rays:- travel in straight lines- are negative in charge- are deflected by electric and magnetic fields- produce sharp shadows- are capable of moving a small paddle wheel
8 Subatomic Particles: Electron This was the discovery of the fundamental unit of charge – the electron.
9 Subatomic Particles: Protons Eugen Goldstein, a German physicist, first observed protons in 1886:Thompson determined the proton’s characteristics.Thompson showed that atoms contained both positive and negative charges.This disproved the Dalton model of the atom which held that atoms were indivisible.But the mass of the atom could not be accounted for by the mass of protons inside it. There had to be something else.
10 James Chadwick discovered the neutron in 1932. Subatomic Particles: NeutronJames Chadwick discovered the neutron in 1932.Its actual mass is slightly greater than the mass of a proton.
11 4.11.00731.0087Memorize name, symbol, relative charge, relative mass, and LOCATION (in nucleus or in orbitals outside of nucleus) of these three subatomic particles. See table 4.1 on page 90/96 in your book.
12 Radioactivity was discovered by Becquerel in 1896. The Nuclear AtomRadioactivity was discovered by Becquerel in 1896.Radioactive elements spontaneously emit alpha (a) particles, beta (b) particles and gamma (g) rays from their nuclei.By 1907, Rutherford found that a particles emitted by certain radioactive elements were helium nuclei, consisting of 2 protons and 2 neutrons.
13 The Rutherford Experiment In 1911, Rutherford performed experiments that shot a stream of a particles at a thin piece of gold foil.Most of the a particles passed through the foil with little or no deflection.He found that a few were deflected at large angles and some a particles even bounced back.
15 The Rutherford Experiment An electron with a mass of amu could not have deflected an a particle with a mass of ~4 amu.Rutherford knew that like charges repel.Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.
16 The Rutherford Experiment If a positive a particle approached close enough to the positive mass it was deflected.Most of the a particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space.Because a particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense.
17 Deflection Scattering Deflection and scattering of a particles by positive gold nuclei.5.5
18 General Arrangement of Subatomic Particles Rutherford’s experiment showed that an atom had a dense, positively charged nucleus.Chadwick’s work in 1932 demonstrated the atom contains neutrons.Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.
19 General Arrangement of Subatomic Particles Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center.The negative electrons surround the nucleus.The nucleus contains protons and neutrons
21 ATOMS AND THEIR SUBATOMIC PARTICLES The atomic number of an element is equal to the number of protons in the nucleus of that element.The atomic number of an atom determines which element the atom is.Every atom with an atomic number of 1 is a hydrogen atom. Every carbon atom contains 6 protons in its nucleus.
22 atomic number1HEvery atom with an atomic number of 1 is a hydrogen atom.1 proton in the nucleus
23 atomic number6CEvery atom with an atomic number of 6 is a carbon atom.6 protons in the nucleus
24 Ions: atoms that have lost or gained electrons Positive ions were explained by assuming that a neutral atom loses electrons.Negative ions were explained by assuming that extra electrons can be added to atoms.
25 When one or more electrons are lost from an atom, a cation is formed. 5.4
26 When one or more electrons are added to a neutral atom, an anion is formed. 5.4
28 Isotopes of the Elements Atoms of the same element can have different masses.They always have the same number of protons, but they can have different numbers of neutrons in their nuclei.The difference in the number of neutrons accounts for the difference in mass.These are isotopes of the same element.
32 Hydrogen has three isotopes 1 proton0 neutrons1 proton1 neutron1 proton2 neutrons
33 Examples of IsotopesElement Protons Electrons Neutrons SymbolHydrogenHydrogenHydrogenUraniumUraniumChlorineChlorine
34 The mass of a single atom is too small to measure on a balance. Atomic MassThe mass of a single atom is too small to measure on a balance.Using a mass spectrometer, the mass of the hydrogen atom was determined.The mass of one hydrogen atom was determined to be x g.
35 A Modern Mass Spectrometer Positive ions formed from sample.Electrical field at slits accelerates positive ions.Deflection of positive ions occurs at magnetic field.From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined.A mass spectrogram is recorded.5.8
36 A typical reading from a mass spectrometer A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given.5.9
37 ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS Single atomic masses are too small to weigh on a balance. To overcome this problem a system of relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers.The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon, carbon-12.A mass of exactly 12 atomic mass units (amu) was assigned to the carbon-12 atom. An amu is defined as exactly equal to 1/12th mass of a carbon-12 atom. 1 amu = x gIsotopes of the same element have different masses.The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12.
38 ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS Calculating relative atomic mass:Convert percent abundance to fraction by dividing by 100Multiply fraction abundance times each isotopes massSum up the resultsRel Atomic Mass = f1*m1 + f2*m2 + …See example on next slide
39 Average atomic mass (amu) To calculate the atomic mass multiply the atomic mass of each isotope by its fractional abundance and add the results.IsotopeIsotopic mass (amu)Abundance (%)Average atomic mass (amu)69.0930.9163.55( amu)=43.48 amu( amu)=20.07 amu63.55 amu
40 Relationship between Mass Number and Atomic Mass Atomic mass found on the Periodic Table is RELATIVE atomic mass, the weighted average of all the isotopes.Each isotope of an element has a unique WHOLE number called the mass number, the sum of protons and neutrons. Mass number is NOT found on most Periodic Tables.******************************************************************ATOMIC MASS IS NOT THE SAME AS MASS NUMBERREMEMBER: ATOMIC MASS IS AN AVERAGE OF THE DIFFERENT ISOTOPES’ MASSES. It will have decimal points!
41 mass number atomic number number of neutrons - = 109 - 47 = 62 Relationship Between Mass Number and Atomic NumberThe mass number minus the atomic number equals the number of neutrons in the nucleus.mass numberatomic numbermass numberatomic numbernumber of neutrons-=109-47=62
42 Atomic Mass, Mass Number & Abundance of Isotopes EXAMPLE: Chlorine has two isotopes, chlorine-35 and chlorine-37. Write down the atomic number, mass number and number of neutrons for each. Then look up the atomic mass. Which isotope is in greater abundance?Cl-35: Z = 17, A = 35, = 18 neutronsCl-37: Z = 17, A = 37, = 20 neutronsAtomic mass is amu.Closer to 35, therefore chlorine-35 is more abundant.
43 Elements are not distributed equally by nature. Oxygen is the most abundant element in the human body (65%).Oxygen is the most abundant element in the crust of the earth (49.2%).In the universe, the most abundant element is hydrogen (91%) and the second most abundant element is helium (8.75%).
44 Distribution of the common elements in nature. 3.2
45 Sources of Element Names Greek-ColorIodine: from the Greek iodes meaning violet.Latin-PropertyFluorine: from the Latin fluere meaning to flow. The fluorine containing ore fluorospar is low melting.German-ColorBismuth: from the German weisse mass which means white mass.LocationGermanium: discovered in 1866 by a German chemist.Famous-ScientistsEinsteinium: named for Albert Einstein.
46 These symbols have carried over from the earlier names of the elements (usually Latin). A number of symbols appear to have no connection with the element.Most symbols start with the same letter as the element.
48 Element Names and Symbols Element names and symbols are in Periodic Table inside front cover.Symbols: Learn first 36 on Periodic Table, plus Ag, Au, Pt, Hg, Sn, Pb, I, U, Ba, and Rn.Also memorize that some elements exist as diatomic molecules : H2, O2, N2, F2, Cl2, Br2, and I2
49 N 7 THE PERIODIC TABLE OF THE ELEMENTS The periodic table was designed by Dimitri Mendelev in 1869.In the table each element’s symbol is placed inside of a box.Above the symbol of the element is its atomic number.7N
50 The elements are arranged in order of increasing atomic number. Elements with similar chemical properties are organized in columns called families or groups .HeNeThese elements are known as the noble gases. They are nonreactive.ArKrXeRn
51 THE PERIODIC TABLE OF THE ELEMENTS 18 columns - called groups7 rows - called periodsRepresentative elements are in groups 1,2, 3A-8A (13-18)Transition metals are in groups 3-12Groups of elements have similar chemical properties
52 THE PERIODIC TABLE OF THE ELEMENTS SPECIAL GROUP NAMES:1A (1) = alkali metals2A (2) = alkaline earth metals8A (18) = noble gases (six)7A (17) = halogens (four)NOTE: H really belongs to its own group, which is why it’s shown by itself in my periodic table! It has 1 electron in outer shell, like the 1A elements, but it’s not a metal, and reacts more like group 7A.
53 The Periodic Table of Elements Nonmetals: upper rightMetals: lower leftMetalloids touch stairs, except for Al.
54 Metals are solid at room temperature. Mercury is an exception. At room temperature it is a liquid.Metals are good conductors of heat and electricity.Most elementsare metalsphysicalproperties of metalsMetals are malleable (they can be rolled or hammered into sheets).Metals have high luster (they are shiny).
55 Metals have high densities Metals are ductile (they can be drawn into wires).Most metals have a high melting point.Most elementsare metalsMetals have high densities
56 Chemical Properties of Metals Metals have little tendency to combine with each other to form compounds.Many metals readily combine with nonmetals to form ionic compounds.They can combine with sulfur.oxygen.chlorine.In nature, minerals are formed by combinations of the more reactive metals with other elements.
57 Chemical Properties of Metals A few of the less reactive metals such as copper, silver and gold are found in the free state.Metals can mix with each other to form alloys.Brass is a mixture of copper and zinc.Bronze is a mixture of copper and tin.Steel is a mixture of carbon and iron.
58 Physical Properties of Nonmetals Everything metals are, nonmetals are NOT!Lack luster (they are dull)Have relatively low melting pointsHave low densities.Poor conductors of heat and electricityAt room temperature, carbon, phosphorous, sulfur, selenium, and iodine are solids.
59 Physical State at Room Temperature phosphorouscarbonSolidsulfurseleniumiodineNONMETALS
60 Physical State at Room Temperature liquidbromineNONMETALS
61 gas nitrogen, oxygen fluorine, chlorine Physical State at Room Temperaturegasnitrogen, oxygenfluorine, chlorinehelium, neon, argon, krypton, xenon, radonNONMETALS
62 The Metalloids boron silicon germanium arsenic antimony tellurium poloniumMetalloids have properties that are intermediate between metals and nonmetals
63 = Alkali Metals = Alkali Earth Metals = Noble Gases = Halogens Chem 10 Spring 2008 Chp 4= Alkali Metals= Alkali Earth Metals= Noble Gases= Halogens= Lanthanides= Actinides= Transition Metalsadd pictures of elements from text
64 Important Groups - Hydrogen NonmetallicColorless, diatomic gas – exists as H2Very low melting point & densityReacts with nonmetals to form molecular compoundsHCl is acidic gasH2O is a liquidReacts with metals to form hydridesmetal hydrides react with water to form H2HX dissolves in water to form acids
65 Important Groups – IA, Alkali Metals Hydrogen usually placed here, though it doesn’t belongSoft, low melting points,low densityFlame tests ® Li = red, Na = yellow, K = violetVery reactive, never find uncombined in natureTend to form water soluble compounds with colorless solutionsReact with water to form basic (alkaline) solutions and H22 Na + 2 H2O ® 2 NaOH + H2releases a lot of heatlithiumsodiumpotassiumrubidiumcesium
66 Important Groups – IIA, Alkali Earth Metals Harder, higher melting, and denser than alkali metalsFlame tests ® Ca = red, Sr = red, Ba = yellow-greenReactive, but less than corresponding alkali metalForm stable, insoluble oxides from which they are normally extractedOxides are basic = alkaline earthReactivity with water to form H2, ® Be = none; Mg = steam; Ca, Sr, Ba = cold waterberylliummagnesiumcalciumstrontiumbarium
67 Important Groups – VIIA, Halogens NonmetalsF2 & Cl2 gases; Br2 liquid; I2 solidAll diatomicVery reactiveCl2, Br2 react slowly with waterBr2 + H2O ® HBr + HOBrReact with metals to form ionic compoundsHX all acidsHF weak < HCl < HBr < HIfluorinechlorinebromineiodine
68 Important Groups – VIIIA, Noble Gases All gases at room temperatureVery low melting and boiling pointsVery unreactive, practically inertVery hard to remove electron from or give an electron to
69 Review What is the atomic number of boron, B? What is the atomic mass of silicon, Si?How many protons does a chlorine atom have?How many electrons does a neutral neon atom have?Will an atom with 6 protons, 6 neutrons and 6 electrons be electrically neutral?Will an atom with 27 protons, 32 neutrons and 27 electrons be electrically neutral?Will a Na atom with 10 electrons be electrically neutral?What is the mass number of fluorine?