1. CHEMISTRY: PACKET #3 Atomic Concepts Atomic Concepts Reference Table: Table S & PT www.regentsprep.org

2. SUBATOMIC PARTICLES ParticleChargeMassLocation Proton+1 1 amu (atomic mass unit) Nucleus Neutron0 1 amu Nucleus Electron 1/1836 or 0 amu Outside Nucleus

3. The Atom Atoms are the building blocks of all matter. Atoms are the building blocks of all matter. The atom is composed of 3 subatomic particles: protons, neutrons, and electrons. The atom is composed of 3 subatomic particles: protons, neutrons, and electrons. Protons and electrons have equal but opposite charges. A neutral atom must contain equal numbers of both. Each atom has a nucleus with an overall positive charge. Within the nucleus are the protons and the neutrons. Surrounding the nucleus are negatively charged electrons. Protons and electrons have equal but opposite charges. A neutral atom must contain equal numbers of both. Each atom has a nucleus with an overall positive charge. Within the nucleus are the protons and the neutrons. Surrounding the nucleus are negatively charged electrons. An atoms identity is defined entirely by the number of protons in the nucleus; the number of protons of any given element NEVER changes. An atoms identity is defined entirely by the number of protons in the nucleus; the number of protons of any given element NEVER changes.

4. AMU (atomic mass unit) The mass of an atom is so extremely small that the atomic mass scale replaces grams as the unit used to describe the masses of atoms The mass of an atom is so extremely small that the atomic mass scale replaces grams as the unit used to describe the masses of atoms The nucleus of a carbon atom containing 6 protons and 6 neutrons is taken as the standard mass for this scale The nucleus of a carbon atom containing 6 protons and 6 neutrons is taken as the standard mass for this scale The amu is defined as 1/12 the mass of a Carbon atom or 1.66x10 -24 g The amu is defined as 1/12 the mass of a Carbon atom or 1.66x10 -24 g

5. The Atomic Model The discovery of the atom as we know it today was a progression, like the discovery of DNA or any other major scientific discovery. Many scientists contributed to the development of present day atomic theory. Each proposed model of the atom was based on the models developed prior to it. With each new discovery dealing with the nature of the atom, a new atomic model was constructed. The discovery of the atom as we know it today was a progression, like the discovery of DNA or any other major scientific discovery. Many scientists contributed to the development of present day atomic theory. Each proposed model of the atom was based on the models developed prior to it. With each new discovery dealing with the nature of the atom, a new atomic model was constructed.

6. HISTORY OF ATOMIC THEORY ATOMIC STRUCTURE DEMOCRITUS(460-370BC)  Thought the material world must be made of atomos -tiny, indivisible, uncuttable particles

7. History of Atomic Theory John Dalton (1766-1844)  Also supported the notion of the atom  He defined an atom as: The smallest particle of an element that retains the chemical identity of that element.  Came up with 4 postulates that outlined his theory about atoms

8. Dalton’s Atomic Theory (1700s) All matter is made up of tiny particles called atoms All matter is made up of tiny particles called atoms All atoms of a given element are alike; while atoms of different elements are different because they have different masses All atoms of a given element are alike; while atoms of different elements are different because they have different masses Compounds are formed when atoms of different elements combine in fixed proportions Compounds are formed when atoms of different elements combine in fixed proportions Atoms are indivisible; atoms can neither be created nor destroyed in chemical reactions. A chemical reaction involves a rearrangement of atoms, not a change in the atoms themselves Atoms are indivisible; atoms can neither be created nor destroyed in chemical reactions. A chemical reaction involves a rearrangement of atoms, not a change in the atoms themselves

9. J.J. Thompson (1890s) Cathode Ray Tube Experiment Cathode Ray Tube Experiment Discovered Electrons, particles with a negative charge. Discovered Electrons, particles with a negative charge. Cathode Ray Tube is a Sealed Glass Tube that contains a gas and has separated metals plates connected to external wires that pass an electrical current through the tube Cathode Ray Tube is a Sealed Glass Tube that contains a gas and has separated metals plates connected to external wires that pass an electrical current through the tube

10. Plum Pudding Model The atom was a hard sphere of positive charge with electrons embedded in it. The atom was a hard sphere of positive charge with electrons embedded in it. Think of it as a chocolate chip cookie; positive- dough & negative- chocolate chips Think of it as a chocolate chip cookie; positive- dough & negative- chocolate chips

11. Further Development of the Atomic Model Ernest Rutherford Ernest Rutherford (1910): disproved Thomson’s model Ernest Rutherford (1910): disproved Thomson’s model Gold Foil Experiment Gold Foil Experiment VERY IMPORTANT! VERY IMPORTANT!

12. Further Development of the Atomic Model Ernest Rutherford Gold Foil Experiment - Observations MOST of the alpha particles (+) went straight through MOST of the alpha particles (+) went straight through SOME of the alpha particles (+) deflected back (bounced back) to the source. SOME of the alpha particles (+) deflected back (bounced back) to the source. Rutherford’s Explanation Most of the α particles pass directly through the gold foil because the gold foil atoms are composed of mostly empty space. Most of the α particles pass directly through the gold foil because the gold foil atoms are composed of mostly empty space. The particles that bounced back did so because they “directly hit” the positive center of the atoms (like charges deflect) The particles that bounced back did so because they “directly hit” the positive center of the atoms (like charges deflect)

13. Bohr Model (1911) Small, dense, positively charged nucleus surrounded by electrons in circular orbits. (a.k.a. the “planetary model”) Small, dense, positively charged nucleus surrounded by electrons in circular orbits. (a.k.a. the “planetary model”) Bohr’s model is fundamentally incorrect because electrons do not move in fixed orbits around the nucleus like the planets do around the sun. Bohr’s model is fundamentally incorrect because electrons do not move in fixed orbits around the nucleus like the planets do around the sun.

14. Bohr’s “Planetary” Atom Model

15. Wave-Mechanical Model (Modern Atomic Theory) Energy and matter are now viewed as acting as both waves and particles Energy and matter are now viewed as acting as both waves and particles Waves are made up of tiny packets called quanta that act like particles Waves are made up of tiny packets called quanta that act like particles Small, dense, positively charged nucleus surrounded by electrons moving in "electron cloud". Small, dense, positively charged nucleus surrounded by electrons moving in "electron cloud". "Orbitals" are areas where an electron with a certain amount of energy is most likely to be found. "Orbitals" are areas where an electron with a certain amount of energy is most likely to be found.

16. Wave-Mechanical Model

17. Orbit vs. Orbital An orbit describes a particular path that an object follows as it travels around another object An orbit describes a particular path that an object follows as it travels around another object For example, the moon has an orbit about the earth For example, the moon has an orbit about the earth Electrons do not follow a particular path around the nucleus Electrons do not follow a particular path around the nucleus Instead, an orbital describes the areas around the nucleus where an electron is most likely to be found (probability of location) Instead, an orbital describes the areas around the nucleus where an electron is most likely to be found (probability of location) The exact path of an electron in this area is not known The exact path of an electron in this area is not known

18. Atomic Theory from Past to Present

20. Atomic Number Located on the lower left hand in the box of the individual element on the Periodic Table. The atomic number is equal to the number of protons (and in a neutral atom, the # of electrons). Located on the lower left hand in the box of the individual element on the Periodic Table. The atomic number is equal to the number of protons (and in a neutral atom, the # of electrons). An element is identified by its number of protons which NEVER CHANGES, and therefore an elements atomic number NEVER CHANGES!! An element is identified by its number of protons which NEVER CHANGES, and therefore an elements atomic number NEVER CHANGES!!

21. Mass Number (Rounded Atomic Mass) Located on the upper left corner in the box of the individual element on the Periodic Table. Located on the upper left corner in the box of the individual element on the Periodic Table. When we talk about the mass number (the “rounded” number), it is the total number of protons and neutrons that an element contains. When we talk about the mass number (the “rounded” number), it is the total number of protons and neutrons that an element contains. In order to calculate the number of neutrons in a given element the formula is In order to calculate the number of neutrons in a given element the formula is Neutrons = Mass # - Atomic #

22. ATOMIC MASS The actual definition of the atomic mass is the average of all of the naturally occurring isotopes of a given element. The actual definition of the atomic mass is the average of all of the naturally occurring isotopes of a given element.

23. ISOTOPES Atoms of the same element that have different numbers of neutrons. Atoms of the same element that have different numbers of neutrons. There are two methods of identifying isotopes. There are two methods of identifying isotopes.

24. Calculating Average Atomic Mass Most elements occur in nature as mixtures of isotopes: Ex) in nature: 11 C, 12 C, 13 C, 14 C Most elements occur in nature as mixtures of isotopes: Ex) in nature: 11 C, 12 C, 13 C, 14 C The average atomic mass of an element can be determined by taking into account the masses of all the naturally occurring isotopes of a particular element found in nature as well as their relative abundances. The average atomic mass of an element can be determined by taking into account the masses of all the naturally occurring isotopes of a particular element found in nature as well as their relative abundances.

25. Calculating Average Atomic Mass Steps 1) Find each isotopes’ fractional abundance (i.e. convert % abundance  decimal) and multiply the fractional abundance of each isotope by their atomic mass. 2) Add up the products Ex 1) Naturally occurring chlorine is 75.78% 35 Cl, which has an atomic mass of 34.969 amu, and 24.22% 37 Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass (that is, the atomic weight) of chlorine.

26. Calculating Average Atomic Mass Ex 2) Three isotopes of silicon occur in nature: 28 Si (92.23%), which has an atomic mass of 27.97693 amu; 29 Si (4.68%), which has an atomic mass of 28.97649 amu; and 30 Si (3.09%), which has an atomic mass of 29.97377 amu. Calculate the atomic weight of silicon.

27. Electron Configurations Bohr’s model describes electrons in terms of energy levels. Electrons that are close to the nucleus have a low energy level and electron farther from the nucleus have a higher energy level. Bohr’s model describes electrons in terms of energy levels. Electrons that are close to the nucleus have a low energy level and electron farther from the nucleus have a higher energy level. Based on Bohr’s model we get Principle Energy Levels. The energy level shows how far the electron is from the nucleus the first energy level is closest to the nucleus and the others are further away. Electrons in the first level have the lowest energy and the energy of the electron increases as the levels increase. Based on Bohr’s model we get Principle Energy Levels. The energy level shows how far the electron is from the nucleus the first energy level is closest to the nucleus and the others are further away. Electrons in the first level have the lowest energy and the energy of the electron increases as the levels increase. Located on the lower left corner, below the atomic number. Located on the lower left corner, below the atomic number.

28. FIRST PRINCIPLE ENERGY LEVEL: holds only 2 electrons. FIRST PRINCIPLE ENERGY LEVEL: holds only 2 electrons. SECOND PRINCIPLE ENERGY LEVEL: holds only 8 electrons. SECOND PRINCIPLE ENERGY LEVEL: holds only 8 electrons. THIRD PRINCIPLE ENERGY LEVEL: holds only 18 electrons. THIRD PRINCIPLE ENERGY LEVEL: holds only 18 electrons. FOURTH PRINCIPLE ENEGY LEVEL: holds only 32 electrons. FOURTH PRINCIPLE ENEGY LEVEL: holds only 32 electrons. Valence Electrons - The number of electrons in the last principle energy level. According to the octet rule, there can be no more than 8 valence electrons. These electrons affect chemical properties of the element. Non-Valence Electrons - All other electrons in an atom other than the last level (valence)

29. Ground State Electrons fill in energy levels and orbitals starting with the one that requires the least energy and progressively move to those levels and orbitals that require increasing amounts of energy. Electrons fill in energy levels and orbitals starting with the one that requires the least energy and progressively move to those levels and orbitals that require increasing amounts of energy. Electrons can gain and lose energy and move to different energy levels. Electrons can gain and lose energy and move to different energy levels. When all electrons are at their lowest possible energy, it is called the "ground state.“ When all electrons are at their lowest possible energy, it is called the "ground state.“ Example: Fluorine (F) has an atomic number of 9, The electron configuration that describes F in the ground state is 2-7

30. Excited State When the electron gains a specific amount of energy from heat, light, or electricity, it moves to a higher orbital and is in the "excited state“ When the electron gains a specific amount of energy from heat, light, or electricity, it moves to a higher orbital and is in the "excited state“ You can use the periodic table to realize an atom is in the excited state You can use the periodic table to realize an atom is in the excited state Look at the electron configuration for sodium (Na) 2-8-1 Look at the electron configuration for sodium (Na) 2-8-1 If Na had the configuration of 2-7-2, you would notice that the second PEL was not completely filled with 8 electrons before the third PEL began to fill. This is an example of an electron configuration in the excited state. Notice that regardless whether it is in the ground or excited state, the total number of electrons never changes. For the example above, the number of electrons remains 11. If Na had the configuration of 2-7-2, you would notice that the second PEL was not completely filled with 8 electrons before the third PEL began to fill. This is an example of an electron configuration in the excited state. Notice that regardless whether it is in the ground or excited state, the total number of electrons never changes. For the example above, the number of electrons remains 11.

31. Flame Test/Bright Line Spectrum In this class we will be performing an experiment called the “Flame Test”. We will be heating up metal powders in order to excite the electrons to jump from a lower PEL to a higher PEL. In this class we will be performing an experiment called the “Flame Test”. We will be heating up metal powders in order to excite the electrons to jump from a lower PEL to a higher PEL. When an electron returns from a higher energy state to a lower energy state, it emits a specific amount of energy usually in the form of light. This is known as a bright line spectrum, and can be used to identify an element like a fingerprint. When an electron returns from a higher energy state to a lower energy state, it emits a specific amount of energy usually in the form of light. This is known as a bright line spectrum, and can be used to identify an element like a fingerprint.

32. Bright Line Spectrum While this light appears as one color to our eyes, it is actually composed of many different wavelengths each of which can be seen using a spectroscope While this light appears as one color to our eyes, it is actually composed of many different wavelengths each of which can be seen using a spectroscope The energy that is given off when an excited electron falls to the ground state is separated into its component wavelengths. The energy that is given off when an excited electron falls to the ground state is separated into its component wavelengths.

33. Each atom has its own distinct pattern of emission lines (bright line spectrum) and these spectra are used to identify elements.

34. Orbitals and Quantum Numbers Bohr’s model introduces two important ideas: 1. Electrons exist only in certain discrete energy levels (quantum numbers) 2. Energy is involved in moving an electron from one level to another. Orbital: A region of most probable location of an electron. Orbital: A region of most probable location of an electron. Quantum Mechanical Model: Does not refer to orbits because the motion and location of an electron around the nucleus cannot be precisely measured. Quantum Mechanical Model: Does not refer to orbits because the motion and location of an electron around the nucleus cannot be precisely measured. Bohr Model Quantum Mechanical Model Bohr Model Quantum Mechanical Model (orbit)(orbital) (orbit)(orbital) Single Quantum # 4 Quantum # ’s Single Quantum # 4 Quantum # ’s

35. The shell with principle quantum # “n” has “n” # of subshells. Ex) n = 1 has 1 subshell: 1s n = 2 has 2 subshells: 2s n = 2 has 2 subshells: 2s 2p 2p n = 3 has 3 subshells: 3s n = 3 has 3 subshells: 3s 3p 3p 3d 3d n = 4 has 4 subshells: 4s n = 4 has 4 subshells: 4s 4p 4p 4d 4d 4f 4f

36. S-orbital S-orbitals are the first orbital to be filled and have the lowest energy. S-orbitals are the first orbital to be filled and have the lowest energy. There is 1 s-orbital per energy level which is capable of holding a one pair of electrons. There is 1 s-orbital per energy level which is capable of holding a one pair of electrons.

37. P-orbitals There are 3 possible p-orbitals, each capable of holding a pair of electrons. There are 3 possible p-orbitals, each capable of holding a pair of electrons. P-orbitals being in the 2 nd energy level. P-orbitals being in the 2 nd energy level.

38. D-orbitals 5 possible d-orbitals, each capable of holding a pair of electrons. 5 possible d-orbitals, each capable of holding a pair of electrons. D-orbitals begin in the 3 rd energy level. D-orbitals begin in the 3 rd energy level. They fill “one energy level late” (in the transition metals). They fill “one energy level late” (in the transition metals).

39. F-orbitals There are 7 possible f-orbitals, each capable of holding a pair of electrons. There are 7 possible f-orbitals, each capable of holding a pair of electrons. F-orbitals fill “2 energy levels behind.” F-orbitals fill “2 energy levels behind.” They begin with the 4 th energy level, but begin filling with electrons after 5d and 6s orbitals. They begin with the 4 th energy level, but begin filling with electrons after 5d and 6s orbitals.

41. Energy Level # Sublevels # Orbitals Max. # Electrons 1 1 (s) 12 2 2 (s, p) 48 3 3 (s, p, d) 918 4 4 (s, p, d, f) 1632

42. REMEMBER! Since any orbital can only hold a maximum of 2 electrons: How many electrons can each subshell hold? Since any orbital can only hold a maximum of 2 electrons: How many electrons can each subshell hold? s subshell can hold a max of _____ electrons s subshell can hold a max of _____ electrons p subshell can hold a max of _____ electrons p subshell can hold a max of _____ electrons d subshell can hold a max of _____ electrons d subshell can hold a max of _____ electrons f subshell can hold a max of _____ electrons f subshell can hold a max of _____ electrons HUND’S RULE: When two or more orbitals are of equal energy, each one is singly occupied before any are doubly occupied.

43. We represent the electron configuration by writing the symbol for the occupied subshell and adding a superscript to indicate the number of electrons in that subshell. We represent the electron configuration by writing the symbol for the occupied subshell and adding a superscript to indicate the number of electrons in that subshell.Ex)

44. Condensed Electron Configurations In writing the condensed electron configuration of an element, the electron configuration of the nearest noble- gas element of lower atomic number is represented by its chemical symbol in brackets. In writing the condensed electron configuration of an element, the electron configuration of the nearest noble- gas element of lower atomic number is represented by its chemical symbol in brackets. This is followed by the full electron configuration for the remaining electrons. This is followed by the full electron configuration for the remaining electrons. Ex) For Na: instead of: 1s 2 2s 2 2p 6 3s 1 Ex) For Na: instead of: 1s 2 2s 2 2p 6 3s 1 we write: [Ne]3s 1 we write: [Ne]3s 1 For Li: instead of 1s 2 2s 1 For Li: instead of 1s 2 2s 1 we write: [He]2s 1 we write: [He]2s 1

45. Exceptions (FYI ONLY!!) Due to the closeness of the 3d and 4f subshell energies: There are a few electron configuration exceptions. Due to the closeness of the 3d and 4f subshell energies: There are a few electron configuration exceptions. Ex) Ex) Cu : [Ar] 4s 1 3d 10 instead of the expected: [Ar] 4s 2 3d 9 Cu : [Ar] 4s 1 3d 10 instead of the expected: [Ar] 4s 2 3d 9 Cr : [Ar] 4s 1 3d 5 instead of the expected: [Ar] 4s 2 3d 4 Cr : [Ar] 4s 1 3d 5 instead of the expected: [Ar] 4s 2 3d 4 Ag: [Kr] 5s 1 4d 10 instead of the expected: [Kr] 5s 2 4d 9 Ag: [Kr] 5s 1 4d 10 instead of the expected: [Kr] 5s 2 4d 9 These “minor departures” are interesting, but not of great chemical significance. These “minor departures” are interesting, but not of great chemical significance.

46. Ions An atom with a charge; a charged particle. An atom with a charge; a charged particle. We can never change the number of protons! So to produce a charge we can add or subtract electrons We can never change the number of protons! So to produce a charge we can add or subtract electrons Atoms lose or gain electrons in order to have a complete outer shell (to become more stable), and follow the octet rule. Atoms lose or gain electrons in order to have a complete outer shell (to become more stable), and follow the octet rule.

47. Example: Sodium Sodium has 11 electrons, with a ground configuration of 2-8-1 Sodium has 11 electrons, with a ground configuration of 2-8-1 The third PEL in incomplete The third PEL in incomplete To become more stable Na will lose the 1 electron in the incomplete PEL. To become more stable Na will lose the 1 electron in the incomplete PEL. When Na loses 1 electron, the overall charge becomes +1, because now Na has 11 protons (never changes), but only 10 electrons. When Na loses 1 electron, the overall charge becomes +1, because now Na has 11 protons (never changes), but only 10 electrons. This ion for lithium is written Na + This ion for lithium is written Na + A positively charged ion is called a cation A positively charged ion is called a cation

48. Na → Na + + e -

49. Example: Chlorine Chlorine has 17 electrons, with a ground configuration of 2-8-7 Chlorine has 17 electrons, with a ground configuration of 2-8-7 The third shell is incomplete (needs 8) The third shell is incomplete (needs 8) To become more stable Cl will gain 1 electron to complete its third shell To become more stable Cl will gain 1 electron to complete its third shell When Cl gains 1 electron, the overall charge becomes - 1, because now Cl has 17 protons (never changes), 18 electrons; - 18 + 17 = - 1 When Cl gains 1 electron, the overall charge becomes - 1, because now Cl has 17 protons (never changes), 18 electrons; - 18 + 17 = - 1 This ion for chlorine is written Cl - This ion for chlorine is written Cl - A negatively charged ion is called an anion. A negatively charged ion is called an anion.

50. Cl + e - → Cl -

51. Lewis Dot Diagrams Helps to show the number of valence electrons in the last principle energy level. There can only be 2 electrons on each side of the symbol of the element. Helps to show the number of valence electrons in the last principle energy level. There can only be 2 electrons on each side of the symbol of the element. 1. Put the 1st and 2nd valence electrons on any side of X. 2. For the 3rd, 4th, and 5th electrons, put each electron on a different side of the symbol. 3. For the 6th, 7th and 8th electrons, add them to any side with 1 electron.

52. REVIEW QUESTIONS 1) The region that is the most probable location of an electron in an atom is A)an orbital B)the nucleus C)an ion D)the excited state 2) Which particles are found in the nucleus of an atom? A) neutrons, only B) protons and electrons C) protons and neutrons D) electrons, only 3) What is the charge of the nucleus in an atom of oxygen-17? A) +17 B) +8 C) 0 D) -2

53. 4) The nucleus of an atom of K-42 contains A) 23 protons and 19 neutrons B) 19 protons and 23 neutrons C) 19 protons and 42 neutrons D) 20 protons and 19 neutrons 5) What is the total number of electrons in a Cu+ ion? A) 29 B) 30 C) 28 D) 36 6) Which Lewis electron-dot structure is drawn correctly for the atom it represents?

54. 7) During a flame test, ions of a specific metal are heated in the flame of a gas burner. A characteristic color of light is emitted by these ions in the flame when the electrons A) gain energy as they return to lower energy levels B) emit energy as they move to higher energy levels C) emit energy as they return to lower energy levels D) gain energy as they move to higher energy levels 8) What is the electron configuration of a sulfur atom in the ground state? A) 2-6 B) 2-4 C) 2-8-6 D) 2-8-4 9) Which electron configuration represents the electrons of an atom in an excited state? A) 2-7-2 B) 2-6 C) 2-8-2 D) 2-4 10. What principle energy level of an atom contains an electron with the lowest energy? A. 1B. 2C. 3D. 4

55. 11) In the early 1900's, evidence was discovered that atoms were not "hard spheres." It was shown that atoms themselves had an internal structure. One experiment involved gold metal foil. (a) In the diagram above, complete the simple model for an atom of gold-197 by placing the correct numbers in the two blanks. (b) In the gold-foil experiment, alpha particles were directed toward the foil. Most of the alpha particles passed directly through the foil with no effect. This result did not agree with the "hard spheres model" for the atom. What conclusion about the internal structure of the atom did this evidence show? (c) In the same experiment, some of the alpha particles returned toward the source. What does this evidence indicate about the charge of the atom's nucleus?