Matter. Forms of Matter

Transcription

1 Matter Everything around us is either matter or energy Matter - everything in universe that occupies space and has mass. Matter composed of tiny particles called atoms and molecules Forms of Matter Matter Pure substance Mixture Compound Element Homogeneous Heterogeneous 1

2 Elements, Atoms and Compounds (oh my!) Elements substance that cannot be separated into simpler substances by chemical means composed of unique kind of atoms carbon (C) different than helium (He) Atom smallest representative particle of an element Compounds* * substance bt composed of two or more elements united chemically in definite proportions. The term molecule used for a specific type of chemical compound held together by covalent bonds Scale of atoms Radius of approximately 1 to 2x10-8 cm!! An angstrom! 1 to 2 x m Feynman s analogy if an apple is magnified to the size of the Earth, then the atoms in the apple are approximately the size of the original i apple Six Easy Pieces 2

3 Mendeleev (1870) Mendeleev summarized behavior of 65 elements When arranged by atomic mass, elements behaved similar characteristics in a periodic manner (periodic law) Most important he predicted properties of elements that had not been discovered! Modern table arranged by atomic number 3

4 Atom composition Composed on nucleus Proton: positively charged subatomic particle in nucleus of atom Neutron : electronically neutral subatomic particle in nucleus of atom Surrounded by electrons Electron: negatively charge subatomic particle in atom with essentially zero mass 4

5 Bohr Model of atom Elements combine to form compounds!! Form compounds through Form compounds through chemical bonds! 5

6 Types of Chemical Bonds Ionic bonds electron transfer Covalent bonds electron sharing Metallic bonds electron pooling II. States of Matter 6

7 Gas What are these states? Liquid Solid Matter has characteristic* physical properties! A physical property is? Examples are Summarize discuss in groups! * Sometimes refer to as observable properties 7

8 Characteristic Physical Properties: Class Summary Physical Property Properties that are measured or observed without changing the composition of the substance Color Taste Odor Texture Density Melting point Boiling point Temperature Hardness Intensive properties [do not depend on the amount of substance] Used to identify substances! 8

9 Physical Property Extensive properties --- depend on quantity of substance [mass, volume, length] Characteristic properties property that is unique to a particular substance Boiling point Melting point Density Solubility Physical Change Substance changes its physical appearance but not its composition Water evaporating Water freezing Still H 2 O! All changes of state are physical changes! 9

10 What is temperature? Temperature is? or Temperature is a physical property that measures? A measure of the average kinetic energy of the particles in a substance What is heat? Energy transferred from a warmer substance to a cooler one based on the temperature difference between the two A perceived warmer hand transfers heat to the colder hand. Cold hands DO NOT transfer COLD to the warmer hands!! 10

11 How do temperature and heat relate to states of matter? Phase Changes and Heat Energy 11

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13 Heating Water Heat added liquid water molecules speed up Water temperature increases Liquid water evaporation increases At boiling point, temperature is constant until all liquid water is converted to vapor Chemical Change 13

14 January 25, Reading: 4.1 to Suggested problems Chapter 3: 49, 57, 63 Chapter 4: 43, 47, 53, 57, 69 The Atom Democritus (c. 460 c. 370 B.C.) proposed that matter was made of discrete indivisible particles. He called his particles atomon, meaning "cannot be cut." His ideas were largely ignored until the scientific revolution of the 16th, 17th, and 18th centuries. 14

15 Lavoisier Law of conservation of mass: The total mass of substances present after a chemical reaction is the same as the total mass of substances before the reaction Proust Law of definite proportions: All samples of a compound have the same composition-the same proportions by mass of the constituent elements. Dalton s Atomic Theory Each element is composed of small particles called atoms. Atoms are neither created nor destroyed in chemical reactions. All atoms of a given element are identical. Compounds are formed when atoms of more than one element combine. 15

16 While atoms were thought to be indivisible, that all changed when J(oseph) J(ohn) Thomson ( ) discovered the electron in Ernest Rutherford ( ) showed that the electrons occupied a region of space surrounding the tiny nucleus. Figuring out just how those electrons behaved required the development of quantum mechanics, a theory in which electrons are treated as wavelike. Cathode ray tube 16

17 Properties of cathode rays Electron m/e = g coulomb -1 Charge on the electron From Robert Millikan showed ionized oil drops can be balanced against the pull of gravity by an electric field. The charge is an integral multiple of the electronic charge, e. 17

18 Radioactivity Radioactivity is the spontaneous emission of radiation from a substance. X-rays and g-rays are high-energy light. -particles are a stream of helium nuclei, He 2+. -particles p l s are a stream of high h speed electrons that originate in the nucleus. The Nuclear Atom Geiger and Rutherford

19 The -particle experiment Most of the mass and all of the positive charge is concentrated in a small region called the nucleus. There are as many electrons outside the nucleus as there are units of positive charge on the nucleus The Nuclear Atom Rutherford protons 1919 James Chadwick neutrons

20 Nuclear Structure Atomic Diameter 10-8 cm Nuclear diameter cm 1 Å Particle Mass Electric Charge kg amu Coulombs (e) Electron Proton Neutron Useful units: Scale of Atoms The heaviest atom has a mass of only g and a diameter of only m. 1 amu (atomic mass unit) = kg 1 pm (picometer) = m 1 Å (Angstrom) = m = 100 pm = cm Biggest atom is 240 amu and is 50 Å across. Typical C-C bond length 154 pm (1.54 Å) Molecular models are 1 Å /inch or about 0.4 Å /cm 20

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22 Main Group elements 22

23 Transition elements Nonmetals 23

24 Metalloids Metals 24

25 Alkali Metals Alkali Earth Metals 25

26 Halogens Nobel gases 26

27 A X Z Chemical symbol A= mass number Z = atomic number A= number of protons + number of neutrons Atomic Mass 27

28 Group 1A lose single s electron in outermost shell. Na Na + + e - Na + has e - configuration of Ne Group 2A lose two s electrons in outermost shell. Ca Ca e - Ca 2+ has e - configuration of Ar Group 7A elements gain 1e -, completing s 2 p 6 in outermost shell Cl + e - Cl - Group 6A elements gain 2e -, completing s 2 p 6 in outermost shell O + 2e - O 2-28

29 EM Radiation Low High Frequency, Wavelength and Velocity Frequency ( ) in Hertz Hz or s -1. Wavelength (λ) in meters m. cm m nm Ă pm (10-2 m) (10-6 m) (10-9 m) (10-10 m) (10-12 m) Velocity (c) ms -1. c = λ λ = c/ = c/λ 29

30 Electromagnetic Spectrum Examples of Interference 30

31 Refraction of Light Slide 10 of 50 General Chemistry: Chapter 8 Prentice-Hall Atomic Spectra (a) (b) (c) (d) (e) Slide 11 of 50 General Chemistry: Chapter 8 Prentice-Hall

32 Atomic Spectra Helium Hydrogen Bohr model, 1 st orbit has radius = 52.9 pm (picometer = m), circumference = nm. Same as just calculated for the e -. When n = 1, one wavelength in standing wave. When n = 2, two wavelengths in standing wave, etc. 32

33 8-4 The Bohr Atom E = -R H n 2 R H = J Energy-Level Diagram ΔE = E f E i = -R H n f 2 1 = R H ( ni 2 n f 2 -R H n i 2 1 ) = h = hc/λ 33

34 Ionization Energy of Hydrogen 1 E = R H ( ni i 2 1 n 2 f ) = h As n f goes to infinity for hydrogen starting in the ground state: h = R H ( ni 2 1 ) = RH This also works for hydrogen-like species such as He + and Li 2+. h = -Z 2 R H Emission and Absorption Spectroscopy 34

35 n=6 n=6 n=5 n=5 h n=4 n=4 n=3 n=3 h h n=2 n=2 n=1 Without the use of a calculator, indicate which of the following transitions in the hydrogen atom results in the emission of light of the greatest energy? 1 1 E 2 2 ni nf 1. n=4 to n=3 2. n=1 to n=2 3. n=3 to n=2 4. n=2 to n=1 5. n=1 to n=3 35

36 Without the use of a calculator, indicate which of the following transitions in the hydrogen atom results in the emission of light of the greatest energy? 1 1 E 2 2 ni nf 1. n=4 to n=3 2. n=1 to n=2 3. n=3 to n=2 4. n=2 to n=1 5. n=1 to n=3 Without the use of a calculator, indicate which of the following transitions in the hydrogen atom results in the emission of light of the longest wavelength? 1 1 E 2 2 ni nf 1. n=4 to n=3 2. n=1 to n=2 3. n=3 to n=2 4. n=2 to n=1 5. n=1 to n=3 36

37 Without the use of a calculator, indicate which of the following transitions in the hydrogen atom results in the emission of light of the longest wavelength? 1 1 E 2 2 ni nf 1. n=4 to n=3 2. n=1 to n=2 3. n=3 to n=2 4. n=2 to n=1 5. n=1 to n=3 Heisenberg s Uncertainty Principle not possible to simultaneously know the exact position and velocity of a particle. Bohr model (planetary model) pictures e - in more detail than possible. de Broglie e - behave as a wave. h n=5 n=4 n=3 n=2 n=1 37

38 Wave motion in restricted systems. Figure 7.13 Only certain allowed wavelengths and energies. Wave must have integral number of wavelengths about circle. If number of waves not integral, wave self-destructs. The best experimental evidence for the existence of discrete (quantized) energy levels in an atom comes from 1. Diffraction of electrons by crystals. 3. Atomic Line Spectra. 2. Dispersion of light by matter. 4. The Stern-Gerlach Experiment 38

39 The best experimental evidence for the existence of discrete (quantized) energy levels in an atom comes from 1. Diffraction of electrons by crystals. 3. Atomic Line Spectra. 2. Dispersion of light by matter. 4. The Stern-Gerlach Experiment Schrödinger equation enables one to calculate allowed energy levels for e -. The allowed energy levels are the same as in the Bohr model. Enables one to calculate the probability of finding the e - Enables one to calculate the probability of finding the e at any particular point in the atom or the amplitude of the wave at any point. 39

40 What is included in the Schrödinger equation? N Potential energy of an electron (attraction between the e - and the nucleus). Kinetic energy of an electron (the energy of the moving e - ). Wave properties of an e -. The solutions to the Schrödinger equation answer the following questions: What is the energy of the electron? Where is the electron most likely to exist? The probability bilit of finding the electron (electron density). 40

41 Electron probability in the ground-state H atom. Higher e - density Bohr radius 0.53Å Figure 7.16 Lower e - density Figures A and B - Electron density (probability of finding e - at a single point) vs. radius for the 1s orbital r R1 s The 1s orbital is represented by drawing a sphere about nucleus within which e - spends 90% of its time, 90% probability contour. 41

42 Figure s 2s orbital is also spherical, but it contains 2 layers. 2s Radial probability plot 1 peak for the 1s, 2 peaks for the 2s. The 3s orbital contains 3 layers. 3s Radial probability plot 3 peaks for the 3s. 42

43 Figure 7.18 The 2p orbitals. 2P z orbital, e - wave has no amplitude on the xy yplane, called a nodal plane Three 2p orbitals shapes. Three 2p orbitals, as far apart as possible Each orbital centered along different axis (x, y, z) P x, P y, P z H atoms, and only H, 2s and 2p have same energy Energy depends ds only on n. In the 3p orbital, the e - spends the most time outside the 2p region, but some time in the 2p region. h hi d h ll h i k h di l In the third shell, there is two peaks on the radial probability plot. 43

44 Figure 7.19 The 3d orbitals. Radial probability plot single peak The lowest energy d orbital is a 3d. Composite of the five 3d orbitals Figure 7.19 continued Five 3d orbitals: 3d xy, 3d xz, 3d yz, 3d x 2 -y2, and 3d z 2 44

45 Electron Configurations and the Periodic Table 45

46 February 6, Reading: 5-1 to Suggested problems Chapter 9: 43, 51, 55, 59, 63, 69, 71, 77, 79, 83, 95 H (Z = 1) Z (atomic number) = number of protons in the nucleus = number of electrons surrounding nucleus in a neutral atom. Quantum number of 1e - : n=1, l=0, m l =0, m s =+1/2 (or 1/2) Electron configuration: (shell) 1s1 (no. of e-) Orbital diagram: one unpaired e - 1s 46

47 He (Z = 2) Two e - s, both n=1, l=0, m l =0, one m s =+1/2, the other 1/2. 1s 2 or 1s Spins paired Li (Z = 3) 1s 2 2s 1 or 1s 2s 2p Why is energy of an e - in 2s < 2p? Determined by effective nuclear charge (Z eff ) that e - experiences. Figure 8.6 Order for filling energy sublevels with electrons Shapes of orbitals 4s orbital is filled before the 3d orbitals. Order of filling can be read directly from the Periodic Table. 47

48 n=2, l=0, m l =0, m s = 1/2. Be (Z = 4) 1s 2 2s 2 or 1s 2s 2p n=2, l=1, m l =-1, m s =+1/2. B (Z = 5) 1s 2 2s 2 2p 1 or 1s 2s 2p Hund s rule all orbitals with the same energy are first filled with 1e -, with all spins in the same direction before any are filled with 2e -. Electrons repel each other. They are as far apart as possible. Pairing of spins is energetically unfavorable. 48

49 n=2, l=1, m l =0, m s =+1/2. C (Z = 6) 1s 2 2s 2 2p 2 or 1s 2s 2p n=2, l=1, m l =+1, m s =+1/2. N (Z = 7) 1s 2 2s 2 2p 3 or 1s 2s 2p n=2, l=1, m l =-1, m s =-1/2. O (Z = 8) 1s 2 2s 2 2p 4 or 1s 2s 2p n=2, l=1, m l =0, m s =-1/2. F (Z = 9) 1s 2 2s 2 2p 5 or 1s 2s 2p 49

50 n=2, l=1,, m l =+1, m s=-1/2. Ne (Z = 10) 1s 2 2s 2 2p 6 or 1s 2s 2p Ne (Z = 10) 1s 2 2s 2 2p 6, 2 nd shell filled Na (Z = 11) 1s 2 2s 2 2p 6 3s 1 or [Ne] 3s 1 Condensed electron configuration [Ne] means all e - in Ne, i.e. a Ne core Ar (Z = 18) [Ne] 3s 2 3p 6 Next fill 4s orbital before 3d orbitals Ca (Z = 20) [Ar] 4s 2 50

51 Transition elements, filling d orbitals 1 st transition series, filling 3d orbitals Sc (Z = 21) [Ar] 4s 2 3d 1 thru Zn (Z = 30) [Ar] 4s 2 3d 10 Two unexpected configurations: Cr (Z = 24) [Ar] 4s 1 3d 5 and Cu (Z = 29) [Ar] 4s 1 3d 10 Half-filled and filled subshells are particularly stable. Periodic table elements arranged by increasing atomic number. Elements with similar properties are in the same group (vertical). Noble gases (8A) and halogens (7A) Devised before any knowledge of electronic structure. Since chemical properties are determined by electronic structure, there is a direct relationship between electronic structure and an elements position on the Periodic Table. 51

52 Figure 8.11 A periodic table of partial ground-state electron configurations In each period, begin filling s orbital in new shell At end of period, have s 2 p 6 in outermost shell 8 e - maximum in outermost shell Period Filling 1 1s 2 2 2s 2 2p 6 3 3s 2 3p 6 4 4s 2 3d 10 4p 6 5 5s 2 4d 10 5p 6 6 6s 2 4f 14 5d 10 6p 6 7 7s 2 5f 14 52

53 Chemical reactivity of atom is determined by its outer electron configuration. These e - are on the surface and contact other atoms. The valence e - are those involved in reactions with other atoms. Main group elements valence e - are s and p e - in outermost shell. Transition elements valence e - are s e - in outermost shell and d e - in n 1 shell. Often if d subshell is filled, these e - are not counted because they don t affect chemical reactivity. For main group elements, no. valence e - = group number Ca (Z = 20) 4s 2 Group 2A Br (Z = 35) 4s 2 4p 5 Group 7A Mn (Z = 25) 4s 2 3d 5 53

54 Going down a group, valence e - stay the same, all have similar chemical properties. In Group 1A, all have ns 1 valence e -. Going across a period, chemical properties of the main group elements change dramatically. Different outer electron configurations. Going across he transition elements, differences are not as dramatic. Differ only in an inner shell. Inner transition elements all have similar properties. 54

55 Figure 8.11 A periodic table of partial ground-state electron configurations Atomic radius think of atoms as spheres Metals - atoms packed as closely as possible. Nonmetals -atoms form covalent bonds by sharing e -. They form individual molecules. Mutual repulsion of their e - defines the closest approach. 55

56 Figure 8.14 Defining metallic and covalent radii Distance between 2 nuclei sum of two atomic radii. Can experimentally measure bond lengths by x-ray crystallography. Cl 2 bond length = 199 pm Cl radius = 100 pm Br 2 bond length = 228 pm Br radius = 114 pm Predict bond length in Cl Br to be 100 pm pm = 214 pm C C distance in diamond 154 pm, C radius is 77 pm. 56

57 On Periodic Table, radius increases going down a group. Within period, radius generally as atomic number. Radius decreases because nuclear charge and Z eff felt by each e - is increasing. Draws e - closer to nucleus, lowers their energy. Figure 8.16 Periodicity of atomic radius 57

58 In 2 nd period, greatest effect from Li (Z = 3) to Be (Z = 4), 152 pm 112 pm Z - eff felt by 2s e almost doubles, from somewhat greater than +1 to somewhat greater than +2. Z increases while inner shielding e - stay constant at 2. Little change from B (Z = 5) to Ne (Z = 10) Contraction from increased nuclear charge counterbalanced by increased e - repulsion. Going down a group, the atomic radius increases. Number of e - shells increases. Nuclear charge also increases, drawing inner shells closer to nucleus, decreasing the energies of these e -. Z eff felt by e - in outermost shell remains relatively constant, t due to shielding by inner e -. 58

59 Compare Ga (Z = 31) with Al (Z = 13). Both are in Group 3A. Coming after the first transition series, gallium has a much larger nuclear charge. Because the 4p electron penetrates the inner shells and experiences a greater Z eff, the radius of gallium is actually smaller than aluminum. Al has a radius of 143 pm Ga has a radius of 135 pm Figure 8.11 A periodic table of partial ground-state electron configurations 59

60 Atomic radii of very heavy atoms not that much larger than light atoms. Li (Z = 3) = 152 pm Cs (Z = 55) = 265 pm First ionization energy (IE 1 ) energy required to remove e - from atom in ground state, while in gas phase. Energy always must be added. Atom (g) + IE 1 Ion + (g) + e - IE 1 < IE 2 < IE 3, etc More difficult to remove e - as + charge of ion increases. 60

61 For Be in Group 2A: Be (900kJ/mol) Be + (1760) Be +2 (14850) Be +3 Why is 3 rd so much higher than 2 nd? e - coming from lower shell. On Periodic Table: IE 1 decreases Radius increases Z eff stays same IE 1 increases Z eff increases Radius decreases Biggest Smallest 61

62 Electron Affinity (EA) energy gained. (EA = +) or released (EA = -) when e - added to atom in gas phase Atom(g) + e - Ion - (g) Most elements have negative EA, i.e. energy released. Going across period, EA tends to become more negative. Halogens (Group 7A) have most negative EA s, form very stable negative ions. 62