Ch. 4 Sec. 1-2, Ch. 3 sec.6-8 ENERGY CHANGES AND THE QUANTUM THEORY THE PERIODIC TABLE

Transcription

1 Ch. 4 Sec. 1-2, Ch. 3 sec.6-8 ENERGY CHANGES AND THE QUANTUM THEORY THE PERIODIC TABLE

2 What Makes Red Light Red? (4.1) Electromagnetic Radiation: energy that travels in waves (light)

3 Waves Amplitude: height of a wave (intensity) Wavelength: distance between crests Frequency: number of wavelengths passing a point in a given amount of time (cycles/ sec., Hz) Speed: light ( 3.00 X 10 8 m/s = c)

4 Waves

5 Waves

6 Waves Higher frequencies correspond to higher energy.

7 Electromagnetic Spectrum

8 Where are Electrons in Atoms? (4.2) Energy levels: fixed areas of energy where the electrons are found (related to distances) Orbitals: regions of space that electrons can occupy

9 Where are Electrons in Atoms? Electrons absorb energy and jump up to higher energy levels. As they return to their original level they give off energy in the form of light. We see different colors depending on the amount of energy released.

10 Atomic Emission Spectrum Atomic emission spectrum: unique pattern of light obtained from an element when subjected to heat or electricity.

11 Atomic Emission Spectrum

12 Quantum leaps Quanta: an energy change of specified size. Ground state: lowest energy state that electrons can occupy. (closest to nucleus) Excited state: a higher energy state that is further from the nucleus. Photon: a particle of light

13 Figure 11.13: Each photon emitted corresponds to a particular energy change.

14 Figure 11.15: The difference between continuous and quantized energy levels.

15 Figure 11.6: Photons of red and blue light.

16 Matter Waves DeBroglie proposed the idea of matter waves. Since light behaves like a wave and particle, matter should also act like a wave and particle. There is a dual nature to matter. Electron microscopes work because matter behaves like a wave.

17 Heisenberg s Uncertainty Principle The position and momentum of a moving object (electron) cannot be measured and known at exactly the same time. We can only tell the probability of finding it in a certain area.

18 IV. New Approach to Atom Quantum-mechanical model uses probability to describe electron motion Electron density cloud that shows the probability of finding an electron in a certain area

19 Figure 11.19: The orbital that describes the hydrogen electron in its lowest possible energy state.

20 Orbitals Area where electron is most likely to be found (90%) s spherical p dumbbell d four leaf clover f - complex

21 Figure 11.20: The hydrogen 1s orbital.

22 Figure 11.25: The three 2p orbitals.

23 Figure 11.28: The shapes and labels of the five 3d orbitals.

24 Orbitals & Energy Principal energy levels described by quantum number (n) Sublevels n=1 1 sublevel 1s n=2 2 sublevels 2s, 2p n=3 3 sublevels 3s, 3p, 3d n=4 4 sublevels 4s, 4p, 4d, 4f 2s is bigger than 1s but same shape

25 Figure 11.24: The relative sizes of the 1s and 2s orbitals of hydrogen.

26 Figure 11.27: Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

27 Orbitals & Energy Sublevel orbitals s 1 p 3 d 5 f 7

28 Electron spin A spinning charge produces a magnetic field, so a spinning electron is like a little magnet. Opposite spins will cancel each other. Only two electrons can fit in an orbital. They are called paired electrons.

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30 Maximum number of electrons in sublevels s 2 p 6 d 10 f 14

31 V. Electron Configurations Show where electrons are located in the atom and how much energy they possess.

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33 The Pauli Exclusion Principle There can only be two electrons per orbital; they must have opposite spins. ( paired & unpaired electrons)

34 Hund s Rule Electrons fill up orbitals in the same sublevel with one electron before they pair up.

35 Orbital diagram

36 Orbital diagram for Oxygen

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38 How are the Elements Organized? (3.6) Periodic law: when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Repeating properties

39 What Kinds Of Chemical Elements Exist? (3.7) Metallic Character: increases from top to bottom down a group and right to left across a period.

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41 Metals, Nonmetals, Semimetals Metals on left side 1. Luster or shine 2. Conduct heat & electricity 3. Usually solids at room temperature 4. Malleable & ductile

42 Metals, Nonmetals, Semimetals Nonmetals on right side 1. No luster 2. Poor conductors of heat & electricity 3. Not malleable or ductile 4. Many are gases at room temperature; some are solids, one is liquid 5. Large variation in properties

43 Metals, Nonmetals, Semimetals Semimetals in between Also known as metalloids Properties are in between metals and nonmetals

44 The Periodic Table? (3.8) In each square 1. element symbol 2. atomic number 3. element s name 4. atomic mass Groups or families: elements in columns; have similar properties Periods: rows

45 Labeling & Naming Groups 3 ways to number (Roman numeral and letter, Arabic numeral and letter, Arabic numbers) Family Names 1A alkali metals 2A alkaline earth metals 7A halogens 8A noble gases

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48 Electron Configurations & Periodic Table Hydrogen 1s 1 Lithium 1s 2 2s 1 Sodium 1s 2 2s 2 2p 6 3s 1 All have one electron in the s orbital at the end. Valence electrons: outermost electrons Elements in groups are similar because they have similar valence electron configurations.

49 Figure 11.31: Orbitals being filled for elements in various parts of the periodic table.

50 D. Abbreviated Electron Configurations Atom s inner (core) electrons: represented by the nearest noble gas with a lower atomic number Li [He]2s 1 Na [Ne]3s 1 K [Ar]4s 1 Rb [Kr]5s 1 Cs [Xe]6s 1

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52 The s, p, d, and f-blocks s-block: first two groups, only holds two electrons p-block: last six groups, can hold 6 electrons d-block: in middle (transition metals), can hold 10 electrons start with 3d f-block: on bottom (inner transition metals), can hold 14 electrons start with 4f

53 Valence Electrons and the Periodic Table Group Number of Valence Electrons 1A one 2A two 3A three 4A four 5A five 6A six 7A seven 8A eight

54 Periodic Trends Atomic radius: size of atom; distance from center (nucleus) to outer electrons 1. Atoms get larger going down a group because of increasing energy level 2. Atoms get smaller from left to right across a period because there are more protons in the nucleus to attract the electrons

55 Figure 11.36: Relative atomic sizes for selected atoms.

56 Periodic Trends Ionization energy: energy needed to remove an electron 1. Decreases down a group 2. Increases across a period (left to right) Small atoms hold electrons more tightly Large atoms lose electrons easily

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58 Periodic Trends Electronegativity: attraction of electrons to atoms involved in a bond Flourine has most (4.0) Cesium & francium least (0.7) 1. Increases from left to right across a period 2. Decreases down a group

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