Standard 1 Atomic Structure Chapters 4-6 vocabulary 1. Atom 2. Atomic number 3. Mass number 4. Atomic mass 5. Isotope 6. Proton 7. Neutron 8. Electron 9. Periodic table 10.Period 11. Group 12.Metal 13. Non-metal 14. Metalloid 15. Alkali metal 16. Alkaline earth metal 17. Halogen 18. Noble gas 19. Ion 20.Valence electron 21. Atomic spectra 22.Atomic orbital 23.Ionization energy 24.Electronegativity Periodic Table. Non-metals Metals Alkali metals Alkaline earth metals Transition metals Semi-metals halogens Nobel gases Metal/non-metal boundary. Summary 1 • Which elements are semi-metals? 1b: groups of the Periodic Table • Metals: – Good conductors – Solid (except mercury) – Lose electrons – Example = aluminum • Non-metals: – poor conductors – Mostly liquid/gas – gain electrons – Example = nitrogen • Semi-metals (metalloids): • Halogens: – Extremely – Have properties of both reactive metals and non-metals – Gain 1 electron – Common use = – Mostly gases – semi-conductors – Example = fluorine – Example = silicon Summary 2 1. Describe the differences between metals and non-metals. 2. Give an example of a metal 3. Give an example of a non-metal 1c: Periodic Groups • Alkali metals – Extremely reactive – Lose 1 electron – Example: sodium • Transition metals – Can lose different numbers of electrons – Example: copper • Alkaline earth metals – Reactive – Lose 2 electrons – Example: calcium • Noble gases – Extremely unreactive – Gases! – Example: helium Summary 3 • Which group of metals are most reactive? 1a: organization of the periodic table • The Periodic Table: organizes elements in groups and periods. • Groups/families: elements have the same physical and chemical properties. • Rows/periods: elements have the same number of electron shells. Summary 4 1. Name another element that would have similar chemical properties to chlorine. 2. Name an atom that is in the same period as chlorine. • The Periodic Table: organizes elements according to atomic number • Atomic number = number of protons Atomic number 6 C 12.011 1 3 2 4 5 6 7 8 9 10 Mass • Mass number: the number of protons and neutrons in an atom (units = amu) • Atomic mass (shown on the periodic table): the average mass of all isotopes • Isotope: an atom with the same number of protons and a different number of neutrons • Note: atomic mass generally increases across the periodic table but not always… (look at atomic number 27&28, 52&53) Isotopes ex: Isotopes of Neon Summary 5 1. What is the mass number for each isotope of neon shown in the example? 2. What is the atomic mass for neon? Standard 1d: electrons • All atoms have an equal number of protons and electrons – Atoms are electrically neutral • Atoms have no charge • Symbol: Ne An equal number of positive protons and negative electrons results in zero charge Summary 6 • How many electrons are in a magnesium atom? • When an atom gains or loses electrons it becomes an ion – Ion = charged particle • number electrons ≠ number protons Na = atom 11 protons 11 electrons Na+ = ion 12 protons 11 electrons Summary 7 • If a magnesium atom loses two electrons, how many electrons will this magnesium ion have? • • • • Valence electrons are: responsible for chemical behavior of atom used for chemical bonding located in the outer orbital Summary 8 1. How many valence electrons does nitrogen have? 2. How many total electrons does nitrogen have? Identifying Atoms by Emission Spectrum: •Adding energy ‘excites’ electrons. •Electrons release energy when they return to the ‘ground state’ (lowest energy level) •Released energy = ‘emission spectrum’ •Each atom has a unique emission spectrum •Scientists use this information in many ways: •CSI can identify elements in an unknown sample •Astronomers can identify elements in stars across the universe Summary 9 What causes an emission spectrum? 1c: Periodic Trends • Electronegativity: The ability of an atom to attract an electron • Example: chlorine is very electronegative because it wants to ______ an electron. • Example: sodium is not very electronegative because it wants to ______ an electron. • General trend for electronegativity: Increasing Increasing electronegativity Note: for noble gases electronegativity = zero Summary 10 1. Which is more electronegative: iodine or chlorine? 2. Which is more electronegative: argon or chlorine? • Ionization energy: the energy needed to remove an electron from an atom • Example: fluorine has a high ionization energy because it wants to ______ an electron. • Example: potassium has a low ionization energy because it wants to ______ an electron. • General trend for ionization energy: Increasing Increasing ionization energy Note: noble gases have a high ionization energy Summary 11 1. Which has a higher ionization energy: iodine or chlorine? 2. Which has a higher ionization energy: argon or chlorine? 3. Which has a lower ionization energy: chlorine or magnesium? • General trend for atomic size (volume) Increasing decreasing Decreasing atomic size Summary 12 • Which is larger: magnesium or calcium? • Which is larger: magnesium or chlorine? General trend for ionic size. • When atoms lose electrons they get much smaller • When atoms gain electrons they get much larger Summary 13 Why is Na+ smaller than Na? Standard 1e: The structure of an atom 1. All the mass of an atom is in the nucleus (Protons & neutrons are in the nucleus) 2. In between the nucleus and the electrons there is only empty space Summary 14 Which particles inside the atom have mass? Earnest Rutherford Rutherford demonstrated that the entire atom is 10,000 times larger than the nucleus • The rutherford experiment: • A stream of positive particles (alpha particles) is aimed at a piece of gold foil. • Only 1 in 8000 particles is deflected (pass close to the gold nucleus). • All other particles travel through ‘empty space’ Summary 15 • How does Rutherford’s experiment demonstrate that an atom is mostly empty space?