Chapter 2
- 1. Topic 2 Atoms, Molecules, and Ions Presented by: Logan
- 2. Dalton’s atomic theory of matter contains four fundamental hypotheses: 1. All matter is composed of tiny indivisible particles called atoms. 2. All atoms of an element are identical in mass and chemical properties, whereas atoms of different elements differ in mass and fundamental chemical properties. 3. A chemical compound is a substance that always contains the same atoms in the same ratio. 4. In chemical reactions, atoms from one or more compounds or elements redistribute or rearrange in relation to other atoms to form one or more new compounds. Atoms themselves do not undergo a change of identity in chemical reactions
- 3. Four minor modifications 1. Not all atoms of an element must have precisely the same mass. 2. Atoms of one element can be transformed into another through nuclear reactions. 3. The compositions of many solid compounds are somewhat variable. 4. Under certain circumstances, some atoms can be divided (split into smaller particles).
- 4. Law of Constant Composition • Also known as the law of definite proportions. • The elemental composition of a pure substance never varies. • The relative amounts of each element in a compound doesn’t vary. H N NH3 ammonia ammonia always has 3 H and 1 N.
- 5. Law of Conservation of Mass The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place. 3H2 + N2 2NH3 ammonia The atoms on the right all appear on the left
- 6. 1. Electrons and protons have electrical charges that are identical in magnitude but opposite in sign. We usually assign relative charges of −1 and +1 to the electron and proton, respectively. 2. Neutrons have approximately the same mass as protons but no charge. They are electrically neutral. 3. The mass of a proton or a neutron is about 1836 times greater than the mass of an electron. Protons and neutrons constitute by far the bulk of the mass of atoms.
- 7. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)
- 8. Isotopes: • Elements are defined by the number of protons. • Atoms of the same element with different masses. • Isotopes have different numbers of neutrons. 11 6C 12 6C 13 6C 14 6C# Neutrons 5 6 7 8
- 9. Average Mass • Because in the real world all the elements exist as mixtures of isotopes. • And we measure many many atoms at a time “Natural abundance” • Average mass is calculated from the isotopes of an element weighted by their relative abundances.
- 10. Average mass, example Isotope abundance Atomic mass 24Mg 78.99% 23.98504 amu 25Mg 10.00% 24.98584 amu 26Mg 11.01% 25.98259 amu Given the above data, what is the average molecular mass of magnesium (Mg)? .7899(23.98504)+0.1000(24.98584)+0.1101(25.98259)= 18.95 + 2.499 + 2.861 = 24.31
- 11. Periodic Table • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties.
- 12. Groups These five groups are known by their names. You gotta know these very well (except for the chalcogens, its far less common).
- 13. Periodic Table Metalloids border the stair-step line (with the exception of Al and Po).
- 14. Elements of life • Elements required for living organisms (pretty much all organisms). • Red, most abundant • blue, next most abundant • Green, trace amounts.
- 15. Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.
- 16. Molecular Compounds Molecular compounds are composed of molecules and almost always contain only nonmetals.
- 17. Diatomic Molecules These seven elements occur naturally as molecules containing two atoms. You should know these guys.
- 18. Types of Formulas • Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. • Molecular formulas give the exact number of atoms of each element in a compound. Example: ethane: Empirical formula: CH3 Molecular formula: C2H6
- 19. Types of Formulas • Structural formulas show the order in which atoms are bonded. • Perspective drawings also show the three-dimensional array of atoms in a compound.
- 20. Ions • When atoms lose or gain electrons, they become ions. Often they lose or gain electrons to have the same number of electrons as the nearest noble gas. • Cations are positive and are formed by elements on the left side of the periodic chart. • Anions are negative and are formed by elements on the right side of the periodic chart.
- 21. Common Cations * * * * * * * * * * * *You should know these.
- 22. Common Anions * * * * * * * * * * * * * * * * * * * * * * * * * *You should know these. ClO2 Chlorite ClO Hypochlorite * *
- 23. More polyatomic anions (the “ites”) SCN- Thiocyanate NO2 - Nitrite HSO3 bisulfite HSO4 bisulfate HPO4 2- Hydrogen phosphate H2PO4 Dihydrogen phosphate ClO- hypochlorite ClO2 - chlorite
- 24. Patterns in Oxyanion Nomenclature • When there are only two oxyanions involving the same element: • The one with fewer oxygens ends in -ite • NO2 − : nitrite; SO3 2− : sulfite • The one with more oxygens ends in -ate • NO3 − : nitrate; SO4 2− : sulfate
- 25. Patterns in Oxyanion Nomenclature • The one with the second fewest oxygens ends in -ite ➢ ClO2 − : chlorite • The one with the second most oxygens ends in -ate ➢ ClO3 − : chlorate • The one with the fewest oxygens has the prefix hypo- and ends in -ite ➢ ClO− : hypochlorite • The one with the most oxygens has the prefix per- and ends in -ate ➢ ClO4 − : perchlorate When there are more than two:
- 26. Examples naming inorganic compounds • Write the name of the cation. • If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. • If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. NaCl sodium chloride NH4NO3 ammonium nitrate Fe(SO4) Iron(II) sulfate KCN potassium cyanide RbOH Rubidium hydroxide LiC2H3O2 lithium acetate NaClO3 sodium chlorate NaClO4 sodium perchlorate K2CrO4 potassium chromate NaH Sodium hydride
- 27. Examples naming inorganic compounds • Write the name of the cation. • If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. • If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. NaCl sodium chloride potasium permanganate KMnO4 Calcium carbonate CaCO3 Calcium bicarbonate Ca(HCO3)2 ammonium dichromate NH4(Cr2O7) potassium phosphate (K)3PO3 Lithium oxide Li2O (O2- is the anion) sodium peroxide Na2O2 (O2 2- is the anion) Calcium sulfide CaS
- 28. Hydrogen • H can be cation or anion • H- hydride • H+ (the cation of an inorganic compound) makes an acid, naming different.
- 29. Acid Nomenclature • If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- : • HCl: hydrochloric acid • HBr: hydrobromic acid • HI: hydroiodic acid
- 30. Acid Nomenclature • If the anion in the acid ends in -ite, change the ending to - ous acid: • HClO: hypochlorous acid • HClO2: chlorous acid
- 31. Acid Nomenclature • If the anion in the acid ends in -ate, change the ending to -ic acid: • HClO3: chloric acid • HClO4: perchloric acid
- 32. Nomenclature of Binary Compounds • The less electronegative atom is usually listed first. • A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.)
- 33. Nomenclature of Binary Compounds • The ending on the more electronegative element is changed to -ide. • CO2: carbon dioxide • CCl4: carbon tetrachloride
- 34. Nomenclature of Binary Compounds If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often merged into one: N2O5: dinitrogen pentoxide not: dinitrogen pentaoxide
- 35. Nomenclature of binary compounds
- 36. Covalent and ionic bonding Electrostatic attraction The bonds in most substances are neither purely ionic nor purely covalent, but they are closer to one of these extremes Ionic compounds consist of positively and negatively charged ions held together by strong electrostatic forces Covalent compounds generally consist of molecules, which are groups of atoms in which one or more pairs of electrons are shared between bonded atoms.
- 37. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.
- 38. Covalent Bonds
- 39. End of Chapter