Alchemy Unit Investigation III. Lesson 7: Life on the Edge

Transcription

1 Alchemy Unit Investigation III Lesson 7: Life on the Edge

2 The Big Question How does the atomic structure of atoms account for the trends in periodicity of the elements?

3 You will be able to: Explain how the Bohr model of an atom can connect the emission spectrum of an atom to the electron structure of the atom.

4 Activity Purpose: Describe the structure of the Bohr model of the atom. (cont.)

5 Notes By using a cathode ray tube, J.J. Thomson discovered that atoms have a small negatively charged particle known as the electron.

6 Notes Later, when James Rutherford fired positively charged alpha particle bullets at a thin gold foil, he found that most the particles passed through the foil. He reasoned that most of the atom must be empty space. Occasionally, one of the particles would be deflected. Rutherford reasoned that the deflection was caused by a small dense core he called the nucleus.

7 Notes Rutherford knew from Thomson s work that an atom has two electrically charged particles a positively charged proton and a small negatively charged electron. When the atoms are found by themselves, they are electrically neutral. This meant that the number of electrons equaled the number of protons.

8 Notes Rutherford proposed that the atom had a small dense positively charged nucleus surrounded by electrons travelling in paths around the nucleus.

9 Notes Since electrons take up the most space and are found on the outside of the atom, can the periodic patterns we see be caused by electrons? For example, what should happen to the size of atoms as they gain more electrons?

10 Notes As you may recall, the sizes of atoms (atomic radii) increases when you go down a period. Elements with atoms that have more electrons seem to have larger atoms.

11 Notes Also remember that When electrons were added to atoms to form anions, they got bigger. When elections were removed from atoms, they became smaller.

12 Notes Electrons may not be the only cause of the periodic patterns. When we travel across a period, the number of electrons increase. But, instead of increasing in size, the size of the atoms DECREASE.

13 Notes We might be able to explain this trend by noting that the number of protons are also increasing. When more protons are present, the nucleus has a greater positive charge to pull the electrons closer.

14 Electrostatic forces in an atom The positively-charged nucleus exerts an attractive force on on the negatively-charged electrons. (red lines) The attractive force is proportional to 1/r, the distance between the electron and nucleus. The loser the electron is to the nucleus, the greater the attractive force.

15 Electrostatic forces in an atom The negatively-charged electrons pushes on on other negatively-charged electrons. (blue lines) This repulsive force is sometimes called screening. Screening reduces the positive force of the nucleus. The screening effect makes the electron shell larger of the atom.

16 Atomic Radii

17 Electron Affinity

18 Ionization Energies

19 Trends in the Periodic Properties One of the first things you should notice about the properties is that the trend or pattern of change across a period repeats itself row after row. Some properties like ionization energies rise to a peak as you go left to right in a row and drops at the beginning of the next row. Some properties like the atomic radius falls as you go left to right in a row and spikes at the beginning of the next row.

20 Problem For example, we can explain the decrease in the atomic radii and the increase in ionization energies across a period by noting that the atom of each successive element has a greater number of protons. That makes the total positive charge increase, pulling the electrons in closer and making them harder to remove.

21 Problem Cations become smaller because the total positive charge of the nucleus stays the same but the electron-electron repulsion decreases. The opposite is true for anions. More electrons means there are more electronelectron repulsions. The problem is explaining the periodic patterns the repeating patterns we see from row to row.

22 Atomic Spectra An explanation for the cause of the periodic trends came from clues scientists found when they were studying the absorption or emission of light-like waves from atom. When atoms of an element are given a large amount of energy, they give off or emit light.

23 Samples of Gas Discharge Tubes

24 Spectra As you can see, the lights from different elements are a different color. When the light from these tubes are passed through a prism, you see a pattern of lines (spectra) unique to each element.

25 Atomic Spectra Electromagnetic radiation are transverse waves that carry energy at the speed of light. Electrons are able to absorb energy from electromagnetic radiation. Electrons give off electromagnetic radiation (light) when they lose energy. Waves with higher frequencies (number of waves per unit time) have shorter wavelengths and carry more energy per unit time than waves with lower frequencies of EM radiation.

26 Electromagnetic Spectrum

27 Gas Discharge Tubes When an electric current is passed through gases in a low pressure tube, the electrons absorb some of the energy. When they lose the extra energy, they release the energy as packets of light or photons. Different gases produce different colors.

28 Copy the following table Gas Color Spectra Observe the gas discharge tubes. Record the color of light produce. Then observe the same light through a prism or diffraction grating.

29 Emission versus Absorption spectra

30 Notice that the emission spectra produces sharp bands of color rather than a full spectrum (rainbow). The production of a band of color means that the electrons are releasing energy (as a color of light) with a single value of energy. The different bands of color suggests that the electrons are able to drop energy by only certain values.

31 Bohr Model of the Atom Niels Bohr developed a planetary model of the atom to explain why excited atoms produced bands of color. Although his model defied a few basic laws of physics, it worked!

32 Bohr s Model Bohr proposed that the electrons traveled in orbits or shells layered around the nucleus, much like the planets in the solar system. The shells would occur at only certain distances from the nucleus

33 Bohr s Model Bohr proposed that Only a fixed number of electrons could go into each shell Electrons in shells farther from the nucleus had greater kinetic energy and were held more loosely than those in the inner shells

34 Bohr Atom

35 Notes Since specific spectral line were made instead of a full rainbow, it was like the electrons were sitting on a shelf and fell to a lower shelf. There was no intermediate height where the electrons could fall from. Bohr imagined that the electrons were moving around on concentric shells around the nucleus.

36 Notes Bohr reasoned that the electrons could absorb energy and move to to higher shell with the additional energy. When the electrons lost energy (conservation of energy), they could drop from an excited state in a higher shell to a lower shell when they gave back the extra energy. The farther the electrons fell, greater the energy it released (blue, purple).

37 Notes These orbits had fixed distances (energy). That meant that the electrons could only absorb certain amounts of energy (only certain wavelength) When the returns to a lower level, it releasea light at only certain wavelengths.

38 Notes In Bohr s model of the atom, the electrons were located in layers around the nucleus. Each layer could hold only a certain number of electrons. In stable atoms, the layers closest to the nucleus were filled before the outer layers could be filled.

39 Notes The first shell or layer could hold a maximum of 2 electrons. The second, 8 electrons The third holds a total of 18 electrons The fourth holds a total 32 The fifth holds a total of 50 # = 2(n 2 ) ; n=1

40 Activity Let s look for periodic patterns in the Bohr model of the atom.

41 Activity Suppose we represent the nucleus by a single sphere Electrons will be replaced by the number and the symbol e - for electrons

42 Activity Here is a section of the periodic table showing the modified representation of Bohr model atoms.

43 Activity How does the number of electrons change as you move from atom to atom across a period or row?

44 Activity How does the number of shells or rings change as you go down a column? Across a period?

45 Activity How do the number of shells in an atom compare to the period number it is found?

46 Activity What happens to the number of electrons and the number of shells when you move from one period to the next? (e.g. neon to sodium)

47 Activity Draw what the unfilled squares would look like.

48 Activity Answers are

49 Notes The outermost shell of the Bohr model of the atom occupied by electrons is called the valence shell. The valence shell contains the valence electrons. (These outer electrons are the ones usually involved in chemical reactions.) All the other electrons (in the inner shells) are called core electrons.

50 Making Sense How many valence electrons do the halogens have? How many do the noble gases have? The Alkali metals?

51 Making Sense What happens to the number of valence electrons as you go down a column (group/family)?

52 Making Sense What happens to the number of valence electrons as you go across period?

53 Making Sense How many core electrons does the element oxygen (O) have? How many does neon (Ne) have? Boron (B)?

54 Activity Recall the periodic table card sort that you did. How were the valence electrons represented on the cards?

55 Notes

56 Notes

57 Notes Bohr s model also provided a possible explanation for the similarity in chemical properties of elements in the same chemical family.

58 Making sense Notice that the elements in the same family have the same number of valence electrons. This suggests that the properties of the elements and their reactivities are related to the number of valence electrons. Each successive period represents a new layer of electrons. The valence electrons are in a new shell that is farther from the nucleus and gets more shielding from the additional electrons in the inner shells.

59 Examples The atomic radius rebounds to a larger size as we move to a new period because the valence electrons are going into a shell that is farther from the nucleus. The atomic radius decreases across a period because the valence electrons see a greater nuclear charge. The ionization energy increases as you go across a period because each additional proton increases the effective nuclear charge to pull the electrons which are in the same shell in closer.

60 Examples The ionization energy decreases as you go to a new period because the new shell is farther from the nucleus. Electronegativity increases as you go across the period because both the effective nuclear charge increases and the distance the shared electrons decrease to increase the attractive force. Cations shed electrons in their valence shell leaving them with only the electrons from the smaller inner core shells.

61 Examples Anions become larger because the number of protons remain the same but the additional electrons increase the number of repulsive interactions. The electron affinity for noble gases suggest that having a filled shell makes an atom very unreactive.