Atomic structure
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1) The document discusses atomic structure and models of the atom including Dalton's model, Rutherford's model, Bohr's model, and quantum numbers. It describes limitations of early models. 2) It also covers types of chemical bonds including ionic bonds formed by electron transfer, covalent bonds formed by electron sharing, and properties of different bond types. 3) Additionally, it discusses properties of different classes of substances like ionic solids, covalent solids, and metals. It explains metallic bonding using the electron sea model.
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Chemistry zimsec chapter 10 chemical periodicity
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Chemistry
Zimsec
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- It discusses the four quantum numbers that describe each electron in an atom: n, l, m, and spin quantum number s.
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2) It describes the different types of bonds - ionic formed between ions with different electronegativities, covalent formed by shared electrons between similar atoms, and metallic formed by delocalized electrons in metals.
3) The properties inferred from different types of bonding are discussed, such as higher melting temperatures for stronger bonds with larger bond energies, and larger coefficients of thermal expansion for weaker bonds.
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Atomic structure
- 2. ATOMIC STRUCTURE Indivisible Of same weight and property in same element Of different weight and property in different elements DEMOCRITUS (5th Century BC) KANAD (6th Century BC) DALTON (1803)
- 3. 1911, Rutherford, Geiger and Marsden 1/20000 α 1897, J J Thomson 0.0004 mm
- 4. Limitation in Rutherford’s model Niels Bohr, 1913
- 5. Max Planck’s Quantum Theory, 1901 Classical worldQuantum world Bohr’s atomic model 1) Privileged Orbitals 2) Constant energy 3) Discrete lines
- 6. (i) Bohr’s Model could not explain the spectra of atoms containing more than one electron. (ii) It could not explain the Zeeman effect. In presence of magnetic field, each spectral line gets split up into fine lines, the phenomenon, is known as Zeeman effect. (iii) It could not explain the Stark effect. In presence of electric field, each spectral line gets split up into fine lines, the phenomenon, is known as Stark effect. Limitation in Bohr’s model
- 8. Address Country State District Post Office Pin Similarly The address of electron is defined by its Quantum Numbers
- 9. First Quantum number (n) This defines the distance of electron from nucleus and hence the shape of orbital is governed by n. Its value may range from 1 to α and this is called Principal Quantum Number (n)
- 10. Second Quantum number is also called the azimuthal quantum number and may have the values from 0 to (n – 1), where n is the principal Quantum number. This governs the shape of the orbital l=0 l=1 l=2 l=3
- 11. ml depends on l. It ranges from – l through 0 to + l i.e it may have 2l+1 values. It defines the orbital orientation in space and is called Third Quantum Number (ml ) Magnetic Quantum Number (ml) px pz py
- 12. Known as Spin Quantum Number Fourth Quantum Number (ms) ms =+1/2 ms =−1/2
- 13. CHEMICAL BONDING • IONIC BONDS • COVALENT BONDS • HYDROGEN BONDS • METALLIC BONDS
- 15. IONIC BOND FORMATION Neutral atoms come near each other. Electron(s) are transferred from the +ve to the -ve atom. They stick together because of electrostatic forces, like magnets.
- 16. IONIC BONDING Metals will tend to lose electrons and become POSITIVE CATIONS Normal sodium atom loses one electron to become sodium ion
- 17. IONIC BONDING Nonmetals will tend to gain electrons and become NEGATIVE ANIONS Normal chlorine atom gains an electron to become a chloride ion
- 19. 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.
- 21. Properties of Ionic Compounds • Crystalline structure. • A regular repeating arrangement of ions in the solid. • Ions are strongly bonded. • Structure is rigid. • High melting points- because of strong forces between ions.
- 23. COVALENT BOND FORMATION When one nonmetal shares one or more electrons with an atom of another nonmetal so both atoms end up with eight valence electrons
- 25. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Atom Oxygen Molecule (O2)
- 26. Linear •Number of Bonds = 2 •Number of Shared Pairs of Electrons = 2 •Bond Angle = 180° EXAMPLE: BeF2
- 27. Bent #1 •Number of Bonds = 2 •Number of Shared Pairs of Electrons = 2 •Number of Unshared Pairs of Electrons = 2 •Bond Angle = < 120° EXAMPLE: H2O
- 28. Tetrahedral •Number of Bonds = 4 •Number of Shared Pairs of Electrons = 4 •Number of Unshared Pairs of Electrons = 0 •Bond Angle = 109.5° EXAMPLE: CH4
- 29. Trigonal Pyramidal •Number of Bonds = 3 •Number of Shared Pairs of Electrons = 4 •Number of Unshared Pairs of Electrons = 1 •Bond Angle = <109.5° EXAMPLE: NH3
- 30. POLAR BOND The more -ve Cl pulls harder on the electrons The electrons spend more time near the Cl H Cl -+ ClH -+
- 31. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.
- 32. Dipole Interactions + - + -
- 33. Hydrogen Bonding H F + + H - H F/O/N + - F H H O H “bond”
- 34. Which substance has the highest boiling point? • HF • NH3 • H2O • WHY? Fluorine has the highest e-neg, SO HF will experience the strongest H bonding and needs the most energy to weaken the i.m.f. and boil
- 35. The Unusual Properties of Water • Unusually high boiling point • Compared to other compounds in Group 16
- 38. Carbon , C Silicon , Si Germanium , Ge [Ar] 3d10 4s2 4p2 [He] 2s2 2p2 [Ne] 3s2 3p2 Allotropes: Diamond, Graphite Silicone Metallic property GeO2 Semiconductor
- 39. METAL 1. Hardness 2. Malleability (Can be Hammered into shape/bend) 3. Ductility (Can be drawn into wires) 4. Elasticity 5. Conductivity Metallic Bond Electron Sea Model