Atomic structure and chemical bonding

1. Atomic structure and chemical bonding

2. Do now • Take out your binder and a pen • Let’s get in the habit of doing this every day

3. Do now: Take out your binder and a pen • Mentally prepare for note taking 

4. Review • What is an atom? An atom is the smallest unit of an element that still retains the properties of that element

5. Structure of an atom • 3 kinds of sub-atomic particles • Protons: positively charged, in nucleus • Neutrons: neutrally charged, in nucleus • Electrons: negatively charged, in cloud around nucleus

6. Atomic Number • The atomic number is the number of protons in an atom. • It is the same for all atoms of a given element. • All hydrogen atoms have one proton • All helium atoms have two protons • And so on…

7. Atomic Number • In an electrically neutral atom (no net charge), the number of protons is equal to the number of electrons. • Therefore the atomic number of an atom is equal to the number of its electrons.

8. Atomic Mass • Protons and neutrons have equal mass • Each has a mass of 1 dalton or 1 atomic mass unit (amu) • Electrons have minimal mass Sooooo….. Atomic mass (in daltons) = # of protons + # of neutrons

9. Practice problem • What is C.C. Carbon’s atomic mass? I have 6 protons, 6 neutrons, and 6 electrons. Do I look fat??!!?!?! Atomic mass = 6 protons + 6 neutrons = 12 daltons *Never ask a lady about her atomic mass.

10. Isotopes • Atoms of the same element with different numbers of neutrons are called isotopes. • Example: Carbon-12 has 6 protons and 6 neutrons • Carbon-13 has 6 protons and 7 neutrons • Carbon-14 has 6 protons and 8 neutrons • They have different atomic masses, but they’re all still carbon. I’m prettier!!! We’re almost identical! Carbon-12 Carbon-14

11. Properties of isotopes • Isotopes of the same element have most of the same properties • They form the same number of chemical bonds • They can substitute for each other in molecules • But they are different in some ways • some isotopes are radioactive, like carbon 14 Carbon-14

12. Radioactive isotopes • What is radioactivity? • Radioactivity is when an unstable nucleus shoots out certain particles in a process called radioactive decay • Alpha particles are made of 2 protons and 2 neutrons • Beta particles are a single electron Carbon-14

13. Real life applications • How is radioactivity useful in my life? • How is it useful in the lives of others? • Page 21 in text • Imaging regions of Alzheimer’s disease • Radiation poisoning at Chernobyl

14. Radiation therapy for thyroid cancer • What is the thyroid? • The thyroid is a gland in your neck • The thyroid gland makes hormones that include iodine atoms • Therefore, iodine is highly concentrated in the thyroid. Normal thyroid Goiter (Iodine deficient thyroid)

15. Radiation therapy for thyroid cancer • Why does this matter for cancer therapy? • Iodine has several isotopes • Iodine-127 is the only stable isotope • Iodine 131 is highly radioactive and emits beta particles • Patients with thyroid cancer are injected with small amounts of iodine-131 • Iodine 131 collects in the thyroid and emits beta particles, which tear through and kill the cancerous tissue.

16. Alzheimer’s treatment • Alzheimer’s is a disease that destroys memory and cognitive skills • Researchers have found high amounts of a protein called beta-amyloidin the brains of Alzheimer’s patients normal brain Alzheimer’s

17. Alzheimer’s treatment • How do we know where beta-amyloid is? • PIB is a protein that binds beta-amyloid • PIB also binds radioactive isotopes • PIB emits beta particles that can be detected • Therefore PIB shows us where beta-amyloid is. normal brain Alzheimer’s

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20. Radioactivity at Chernobyl • Nuclear power plants use radioactive isotopes as fuel. • Radioactive decay events give off huge amounts of heat that is trapped and used for power

21. Radioactivity at Chernobyl • 1986 – nuclear meltdown in Chernobyl, Ukraine • Meltdown sent radioactive isotopes into the air, water, and land • Death from radiation poisoning • Genetic mutations

22. Electrons • “Of the three subatomic particles – protons, neutrons, and electrons – only electrons are directly involved in the chemical activity of an atom.” Why? Because electrons are responsible for forming chemical bonds.

23. Electron Shells • What is an electron shell? • Electrons move around the nucleus only at certain energy levels, called electron shells • The first shell holds two electrons • The second and third shells each hold eight electrons. • Electrons fill up a lower shell first before filling up a higher shell

24. Electron shells and the periodic table • Each row of the periodic table represents one electron shell • The first shell can only hold 2 electrons, so the first row of the periodic table only has 2 elements (hydrogen and helium) • Shells 2 and 3 have 8 electrons. Rows 2 and 3 have 8 elements • Each time you go across a row of the periodic table, you fill in a new shell of electrons

25. Valence shell • The outermost electron shell is called the valence shell • Only electrons in the valence shell form bonds with other atoms. • Atoms form bonds to achieve a full valence shell

26. Practice problem • With a partner, determine how many valence electrons each of the following atoms has: • Hydrogen • Helium • Lithum • Carbon • Nitrogen • Oxygen

27. What if the valence shell is already full? • The right-most column of the periodic table contains the noble gases. • The noble gases are at the very end of their row, so their entire valence shell is full. • Consequently, noble gases do not need to form bonds or take part in reactions

28. Chemical bonding • Atoms with incomplete valence shells interact with other atoms to form full valence shells • These interactions are called chemical bonding • Atoms can share, donate and receive electrons to get to a full valence shell.

29. Covalent bonds • Vocabulary: covalent Break it down! • co = together (like co-worker, cooperate) • valent = valence Covalent = valence electrons shared together

30. Do Now • With a partner, answer the following: • What is a covalent bond?

31. Covalent bonds • Bonds in which two atoms share one or more pairs of valence electrons. • Two or more atoms held together by covalent bonds form a molecule. H H H2

32. Double and Triple bonds • Two atoms can share more than one covalent bond • In a double bond, four electrons are shared • In a triple bond, six electrons are shared Triple bond (N2)

33. Bonding Capacity • In each covalent bond, an atom shares a valence electron and gets a new one shared with it. • Atoms need to have at least one of their own valence electrons for every covalent bond they form. Single bond

34. Practice problem • How many valence electrons does boron have? • What’s the maximum number of covalent bonds boron can form? • Can boron ever achieve a valence octet by covalent bonding alone?

35. Bonding capacity • For each covalent bond an atom forms, it must have a space in its valence shell for a new electron. • Ex: fluorine has 7 electrons in its valence shell. It is only one electron short of a full octet, so it can only form one bond.

36. Bonding capacity • The bottom line: • For each covalent bond an atom forms, it must have: 1) A valence electron to share and 2) a vacancy in its valence shell to accept a new electron.

37. Practice problem • Fill in this table in your notes: Water Ammonia Methane

39. Electronegativity • Not all covalent bonds share electrons equally. • The tendency of an atom to pull electrons toward itself is called electronegativity.

40. Electronegativity • How do we measure electronegativity? • Pauling electronegativity scale • Each element is scored from .7 to 4 (no units) • Higher score = stronger pull on electrons

41. Pauling electronegativity scale

42. Electronegativity trend • Electronegativity increases as: • You move right across a period (row) • Closer to a full octet = less likely to give up electrons • You move up a group (column) • Less electron shells shielding the nucleus = more pull from positive nucleus holding on to electrons

43. Extension question • Explain two reasons why Francium (group 1, period 7) has the lowestelectronegativity of any known element

44. Polar covalent bonds • Polar covalent bonds are bonds in which bonding electrons are not evenly shared • PCBs occur when electronegativities differ by more than .4 and less than 1.7 • This results in a partial negative charge (δ-) on the more electronegative atom and a partial positive charge (δ+) on the less electro- negative atom

45. Electrostatic potential map • Map which shows the density of electrons around regions of a molecule • Electrons are not evenly distributed in polar molecules • More electronegative atoms have higher electron density (red on map) Electronegativities Oxygen = 3.4 Hydrogen = 2.2 Water

46. Challenge question • Why do polar covalent bonds only leave partial charges on participating atoms? • Answer: Because the electrons are still shared, they are just pulled more strongly toward the electronegative atom.

47. Non-polar covalent bonds • When electronegativities differ by less than .4, a covalent bond is considered non-polar. • Electrons are shared equally in these bonds Methane (CH4) Electronegativities Carbon = 2.55 Hydrogen = 2.2

48. Review question • Does the cartoon below show a polar covalent or a non-polar covalent bond? Explain.

49. Ionic bonds • In ionic bonds, electrons are transferred from one atom to another. • Lose e- + charge, gain e- - charge. • Charged atoms are called ions • After e- transfer, the ions attract each other

50. Ionic bonds • What atoms form ionic bonds? • Ionic bonds occur between atoms that differ in electronegativity by more than 1.7 Ion forming atoms: