Atoms, Ions, and the Periodic Table

1. Atoms, Ions, and the Periodic Table What is an atom? It is smallest particle of an element that retains the elements properties. But how did we come to know all the information we have about these tiny particle?

2. Democritus (460-370 BC) • Matter is made of tiny, solid, indivisible particles which he called atoms (from atomos, the Greek word for indivisible). • Different kinds of atoms have different sizes and shapes. • Different properties of matter are due to the differences in size, shape, and movement of atoms. • Democritus’ ideas, though correct, were widely rejected by his peers, most notably Aristotle (384-322 BC). Aristotle was a very influential Greek philosopher who had a different view of matter. He believed that everything was composed of the four elements earth, air, fire, and water. Because at that time in history, Democritus’ ideas about the atom could not be tested experimentally, the opinions of well-known Aristotle won out. Democritus’ ideas were not revived until John Dalton developed his atomic theory in the 19th century!

3. John Dalton (1766-1844) • All matter is composed of extremely small particles called atoms. • All atoms of one element are identical. • Atoms of a given element are different from those of any other element. • Atoms of one element combine with atoms of another element to form compounds. • Atoms are indivisible. In addition, they cannot be created or destroyed, just rearranged.

4. Dalton’s theory was of critical importance. He was able to support his ideas through experimentation, and his work revolutionized scientists’ concept of matter and its smallest building block, the atom. • Dalton’s theory has two flaws: • In point #2, this is not completely true. Isotopes of a given element are not totally identical; they differ in the number of neutrons. Scientists did not at this time know about isotopes. • In point #5, atoms are not indivisible. Atoms are made of even smaller particles (protons, neutrons, electrons). Atoms can be broken down, but only in a nuclear reaction, which Dalton was unfamiliar with.

5. Discovery of the Electron JJ Thomson (1856-1940) • Discovered the electron, and determined that it had a negative charge, by experimentation with cathode ray tubes. A cathode ray tube is a glass tube in which electrons flow due to opposing charges at each end. Televisions and computer monitors contain cathode ray tubes. • Thomson developed a model of the atom called the plum pudding model. It showed evenly distributed negative electrons in a uniform positive cage. • Diagram of plum pudding model:

6. Discovery of the NucleusErnest Rutherford (1871-1937) • Discovered the nucleusof the atom in his famous Gold Foil Experiment. • Alpha particles (helium nuclei) produced from the radioactive decay of polonium streamed toward a sheet of gold foil. To Rutherford’s great surprise, some of the alpha particles bounced off of the gold foil. This meant that they were hitting a dense, relatively large object, which Rutherford called the nucleus.

7. Rutherford then discovered the proton, and next, working with a colleague, James Chadwick (1891-1974), he discovered the neutron as well.

8. Models of the Atom - Niehls Bohr • Developed the Bohr model of the atom (1913) in which electrons are restricted to specific energies and follow paths called orbits a fixed distance from the nucleus. This is similar to the way the planets orbit the sun. However, electrons do not have neat orbits like the planets. • Diagram of Bohr model:

9. Quantum Mechanical Model • This is the current model of the atom. We now know that electrons exist in regions of space around the nucleus, but their paths cannot be predicted. The electron’s motion is random and we can only talk about the probability of an electron being in a certain region.

10. Sub-Atomic ParticlesEach atom contains different numbers of each of the three SUBatomic particles “A neutron walked into a bar and asked how much for a drink. The bartender replied, “For you, no charge.”

11. Atomic Number The periodic table is organized in order of increasing atomic number. The atomic number is the whole number that is unique for each element on the periodic table. The atomic number defines the element. For example, if the atomic number is 6, the element is carbon. If the atomic number is not 6, the element is not carbon. The atomic number represents: • the number of protons in one atom of that element • the number of electrons in one atom of that element (with an ion, the electrons will be different) **Therefore, protons = electrons in a neutral atom**

12. Atomic Mass • mass of an element measured in amu (atomic mass units) • all compared to C-12 (the mass of carbon 12, which has a mass of exactly 12 amu • listed on the periodic table • Mass number= #of protons + # of neutrons

13. Isotopes • Isotopes are atoms of an element with the same number of protons but different numbers of neutrons. • Most elements on the periodic table have more than one naturally occurring isotope. • There are a couple of ways to represent the different isotopes. One way is to put the mass after the name or symbol: Carbon-12 or C-12 • Another way is to write the symbol with both the mass number and atomic number represented in front of the symbol:

14. Determining Average Atomic Mass • The atomic mass on the periodic table is determined using a weighted average of all the isotopes of that atom. • In order to determine the average atomic mass, you convert the percent abundance to a decimal and multiply it by the mass of that isotope. The values for all the isotopes are added to together to get the average atomic mass.

15. Example of Average atomic mass calculation Given: 12C = 98.89% at 12 amu 13C = 1.11% at 13.0034 amu Calculation: (98.89%)(12 amu) + (1.11%)(13.0034 amu) = (0.9889)(12 amu) + (0.011)(13.0034 amu) = 12.01 amu

16. Now you try one: • Neon has 3 isotopes:  Neon-20 has a mass of 19.992 amu and an abundance of 90.51%.  Neon-21 has a mass of 20.994 amu and an abundance of 0.27%.  Neon-22 has a mass of 21.991 amu and an abundance of 9.22%.  What is the average atomic mass of neon? • The answer is: (0.9051)(19.992 amu) + (0.0027)(20.994 amu) + (0.0922)(21.991 amu) = 20.179 amu Now compare this mass for Neon to the mass on the periodic table!

17. Electromagnetic Radiation • Electromagnetic radiation is a form of energy that travels through space in a wave-like pattern. eg. Visible light • It travels in photons, which are tiny particles of energy that travel in a wave like pattern. Although we call them particles, they have no mass. Each photon carries one quantum of energy. • These photons of energy travel at the speed of light (c) = 3.00 x 108 m/s in a vacuum

18. What is a wave and how do we measure it? • Frequency (ν) – number of waves that passes a given point per second (measured in Hz) • Wavelength (λ) – shortest distance between two equivalent points on a wave (measured in m, cm, nm)

19. Electromagnetic spectrum (EM) • The electromagnetic spectrum shows all wavelengths of electromagnetic radiation – the differences in wavelength, energy and frequency differentiates the different types of radiation. • Note that as the wavelength increases, the energy and the frequency decrease.

20. Ground state vs. Excited state • Electrons generally exist in the lowest energy state they can. We call this the ground state. • However, if energy is applied to the electrons, they can be “excited” to a higher energy and we call this an excited state. • The excited state electron doesn’t stay “excited”. It will fall back to the ground state quickly. When the electron returns to the ground state, energy is released in the form of light. One example of this is lasers.

21. Electrons in Atoms • We are most concerned with electrons because electrons are the part of the atom involved in chemical reactions. • Electrons are found outside the nucleus, in a region of space called the electron cloud. • Electrons are organized in energy levels of positive integer value (n = 1, 2, 3,...). • Within each energy level are energy sublevels, designated by a letter: s, p, d, or f. • Each sublevel corresponds to a certain electron cloud shape, called an atomic orbital.

22. The electron cloud is like an apartment building.

23. The energy levels are like floors in the apartment building.

24. The sublevels are like apartments on a floor of the building. Just like there are different sizes of sublevels, there are different sizes of apartments: 1 bedroom, 2 bedroom, etc. The orbitals are like rooms within an apartment.

25. The electrons are like people living in the rooms.

26. What do these orbitals look like? • The s, p, d and f orbitals look different and increase in complexity (f-orbitals not shown… they are very complex)

28. Number of electrons in each sublevel depends on number of orbitals! • Each orbital can hold a maximum of 2 electrons. • An “s” sublevel contains 1 s orbital. How many total electrons can fit in an s sublevel? • 2 • A “p” sublevel contains 3 p orbitals. How many total electrons can fit in a p sublevel? • 6 • A “d” sublevel contains 5 d orbitals. How many total electrons can fit in a d sublevel? • 10 • An “f” sublevel contains 7 f orbitals. How many total electrons can fit in an f sublevel? • 14

29. The Aufbau Principle • Three rules govern the filling of atomic orbitals. The first is: • The Aufbau Principle: Electrons enter orbitals of lowest energy first. The Aufbau order lists the orbitals from lowest to highest energy: (“Aufbau” is from the German verb aufbauen: to build up) 1s2 2s2 2p6 3s2 3p64s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10

30. The Pauli Exclusion Principle • An atomic orbital may hold at most 2 electrons, and they must have opposite spins (called paired spins). • When we draw electrons to show these opposite spin pairs, we represent them with arrows drawn in opposite directions. Write this down in your notes if you haven’t!

31. Write this down in your notes if you haven’t! Hund’s Rule • When electrons occupy orbitals of equal energy (such as three p orbitals), one electron enters each orbital until all the orbitals contain one electron with spins parallel (arrows pointing in the same direction). Second electrons then add to each orbital so that their spins are paired (opposite) with the first electron in the orbital.

32. An electron configuration uses the Aufbau order to show how electrons are distributed within the atomic orbitals. • How to read a segment of an electron configuration: Example 3p6 3 = energy level p = sublevel 6 = # of electrons Now, let’s look at how to put these together for a specific element!

33. Electron Configurations • This is one way to represent the electrons of an atom. We will try a few together: 1s2 2s2 2p2 9 1s2 2s2 2p5 1s2 2s2 2p6 3s2 12 1s2 2s2 2p6 3s2 3p6 18

34. Orbital Diagrams • Orbital diagrams show with arrow notation how the electrons are arranged in atomic orbitals for a given element. ↑↓ ↑↓ ↑ ↑ . 1s 2s 2p 6 ↑↓ ↑↓ ↑↓ ↑↓ ↑ . 1s 2s 2p 9 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓. 1s 2s 2p 3s 12 18 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1s 2s 2p 3s 3p

35. Valence electrons • Electrons in the outer energy level of an atom. They are like the front lines of an army, because they are the ones involved in chemical reactions (valence electrons get shared or transferred during reactions). • The number of valence electrons that an atom has is directly responsible for the atom’s chemical behavior and reactivity. • We can represent the number of valence electrons pictorially by drawing the electrons around the symbol in a “dot diagram”. The electrons are drawn in on each side of the symbol and are not paired up until they need to be. • Eg. . Be.

36. 1s2 2s1 1 Li. 1s2 2s2 2 . Be . 1s2 2s2 2p1 3 .B . ̇ . .C . ̇ 1s2 2s2 2p2 4 . .N : ̇ 5 1s2 2s2 2p3 . :O : ̇ 1s2 2s2 2p4 6 .. :F : ̇ 7 1s2 2s2 2p5 .. :Ne : ̇̇̇̇ 1s2 2s2 2p2 8

37. The Periodic Table • The rows on the periodic table are called periods • The columns on the periodic table are called groups or families • Elements within a group or a family have similar reactivity. What do you know about all elements in a period that could explain this? • They have the same number of valence electrons

38. Since many of the families on the periodic table have such similar properties, they some have specific names that you need to know. Get out your periodic table and label each section as we look at them together.

39. Alkali Metals are group 1 and are the most reactive metals. They form +1 ions by losing their highest energy s1 electron. 1 valence electron. Alkaline Earth Metals are in group 2. the form 2+ ions by losing both of the electrons in the highest energy s orbital. 2 valence electrons. The transition metals include groups 3 through 12 and these metals all lose electrons to form compounds Halogens are in group 17 and they are the most reactive nonmetals. The form -1 ions by gaining 1 electron to fill the highest energy p orbital. They have 7 valence electrons. Noble Gases are in group 18. They do not form ions because they have a full outer shell of electrons and do not need any more electrons. They do not form compounds.8 valence electrons

40. Electromagnetic Radiation • Electromagnetic radiation is a form of energy that travels through space in a wave-like pattern. eg. Visible light • It travels in photons, which are tiny particles of energy that travel in a wave like pattern. Although we call them particles, they have no mass. Each photon carries one quantum of energy. • These photons of energy travel at the speed of light (c) = 3.00 x 108 m/s in a vacuum

41. What is a wave and how do we measure it? • Frequency (ν) – number of waves that passes a given point per second (measured in Hz) • Wavelength (λ) – shortest distance between two equivalent points on a wave (measured in m, cm, nm)

42. Electromagnetic spectrum (EM) • The electromagnetic spectrum shows all wavelengths of electromagnetic radiation – the differences in wavelength, energy and frequency differentiates the different types of radiation. • Note that as the wavelength increases, the energy and the frequency decrease.

43. Ground state vs. Excited state • Electrons generally exist in the lowest energy state they can. We call this the ground state. • However, if energy is applied to the electrons, they can be “excited” to a higher energy and we call this an excited state. • The excited state electron doesn’t stay “excited”. It will fall back to the ground state quickly. When the electron returns to the ground state, energy is released in the form of light. One example of this is lasers.

44. Nuclear Forces • The force that holds the protons together within the nucleus even though there are repulsive forces that would otherwise push the positive protons away from each other. (also known as strong force)

45. Radiation • Radiation-it’s the transfer of energy • Radioactivity-The spontaneous emission of radiation by an unstable nucleus.

46. Good vs. Bad • Ionizing • Has enough energy to kick off an ion. • Very high energy • Non ionizing • Does not have enough energy to kick off an ion • Low energy

47. A. Types of Radiation • Alpha particle () • helium nucleus paper 2+ • Beta particle (-) • electron 1- cardboard • Positron (+) • +’ly charged e- 1+ concrete • Gamma () • high-energy photon thick lead 0

48. parent nuclide alpha particle daughter nuclide B. Nuclear Decay • Alpha Emission Numbers must balance!!

49. B. Nuclear Decay electron • Beta Emission

50. B. Nuclear Decay • Gamma Emission • Usually follows other types of decay. • Transmutation • One element becomes another.