1 / 62

Topic 2 Atoms, Elements, Molecules, Ions, and Compounds

Topic 2 Atoms, Elements, Molecules, Ions, and Compounds.

ianthe
Download Presentation

Topic 2 Atoms, Elements, Molecules, Ions, and Compounds

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Topic 2 Atoms, Elements, Molecules, Ions, and Compounds Early in the 19th century John Dalton developed atomic theory. His theory explained the best available experimental data at that time. His theory has been modified since then with the discovery of other data, but his work was the initial ground work that we will examine first.

  2. Atomic Theory of Matter Protons, neutrons, electrons Later found indivisible to be untrue. • Postulates of Dalton’s Atomic Theory 1.)All matter is composed of indivisible atoms. An atom is an extremely small particle of matter that retains its identity during chemical reactions. 2.)An element is a type of matter composed of only one kind of atom, each atom of a given element having the same properties. Mass is one such property. Thus the atoms of a given element have a characteristic mass. Later found all atoms of the same element does not have to have the same mass. Atoms have different isotopes that have the same # protons but different # neutrons and hence different mass. Note: #protons gives identity of atom.

  3. Atomic Theory of Matter • Postulates of Dalton’s Atomic Theory • 3.) A compound is a type of matter composed of atoms of two or more elements chemically combined in fixed proportions. • The relative numbers of any two kinds of atoms in a compound occur in simple ratios. • Water, for example, consists of hydrogen and oxygen in a 2 to 1 ratio (2H: 1O) for all molecules of water.

  4. Atomic Theory of Matter i.e., solid sodium mixed with chlorine gas forms a new substance, salt, with totally different properties from the starting materials. • Postulates of Dalton’s Atomic Theory • 4.) A chemical reaction consists of the rearrangements of the atoms present in the reacting substances to give new chemical combinations present in the substances formed by the reaction (new chemical with different properties). • 5.) Atoms are not created, destroyed, or broken into smaller particles by any chemical reaction. Na(s) + Cl2(g) 2 NaCl (s) 2 Once again, later found indivisible to be untrue. Protons, neutrons, electrons

  5. Atomic Theory of Matter • The Structure of the Atom • Although Dalton postulated that atoms were indivisible, experiments at the beginning of the 1900’s showed that atoms themselves consist of particles. • Experiments by Ernest Rutherford in 1910 showed that the atom was mostly “empty space.”

  6. Atomic Theory of Matter • These experiments showed that the atom consists of two kinds of particles: a nucleus, the atom’s central core, which is positively charged and contains most of the atom’s mass, andone or more electrons. • Electrons are very light, negatively charged particles that exist in the region around the atom’s positively charged nucleus. Nucleus (+) e-

  7. Atomic Theory of Matter • In 1897, the British physicist J. J. Thompson conducted a series of experiments that showed that atoms were not indivisible particles but instead made of smaller particles. • From his experiments, Thompson calculated the ratio of the electron’s mass, me, to its electric charge, e.

  8. Atomic Theory of Matter • In 1909, U.S. physicist, Robert Millikan obtained the charge on the electron (1.602 x 10-19 C). • These two discoveries (Millikan and Thompson) combined provided us with the electron’s mass of 9.109 x 10-31 kg, which is more than 1800 times smaller than the mass of the lightest atom (hydrogen) thereby proving that the atom is made up of smaller particles. • These experiments showed that the electron was indeed a subatomic particle.

  9. Atomic Theory of Matter The nuclear model of the atom. • Ernest Rutherford, a British physicist, put forth the idea of the nuclear model of the atom in 1911, based on experiments done in his laboratory by Hans Geiger and Ernest Morrison. • Rutherford’s famous gold leaf experiment gave credibility to the theory that the majority of the mass of the atom was concentrated in a very small nucleus. • Positively charged alpha particles were directed at a metal foil. Only 1/8000 were deflected indicating that the nucleus was extremely small and positively charged. Only those alpha particles that directly hit the nucleus were deflected; the rest passed through.

  10. Atomic Theory of Matter • Most of the mass of an atom is in the nucleus; however, the nucleus occupies only a very small portion of the space in the atom. • The diameter of an atom is approximately 100 pm while the diameter of the nucleus is approximately 0.001 pm. For comparison, if an atom was 3 miles in diameter, the nucleus would be the size of a golf ball. • The nucleus of an atom is composed of two different kinds of particles: protons (+)and neutrons (neutral). • An important property of the nucleus is its positive electric charge.

  11. Atomic Theory of Matter • A proton is the nuclear particle having a positive charge equal to that of the electron’s (a “unit” charge) and a mass more than 1800 times that of the electron. It is for this reason that we refer to H as a pure proton. • The number of protons in the nucleus of an atom is referred to as its atomic number (Z) and gives the identity of an element. All species that have same #p have the same properties. H Z=1 1p, 1e- neutral species: #p = #e- Na Z=11 11p, 11e- Cl Z=17 17p, 17e- Na+ Z=11 11p, 10e- Cl- Z=17 17p, 18e- + charge, more p than e- - charge, more e- than p #p remains constant in species; #e- can vary and dictates the charge of species

  12. Atomic Theory of Matter • An element is a substance whose atoms all have the same atomic number (Z). The #protons defines the identity of an atom and can be found on the periodic table (large number in top of element box). • The neutron is a nuclear particle having a mass almost identical to that of a proton, but no electric charge. The charge of the nucleus comes from the #protons. The atoms may have different masses because of different #neutrons (isotopes). • Summary of masses and charges of the three fundamental particles:

  13. The mass number (A) is the total number of protons and neutronsin a nucleus. A = #p + #n = Z + #n How many neutrons does sodium 23 have? A = 23, Z = 11(number on periodic table) #n = 23 - 11 = 12 A = Z + #n 23 = 11 + #n • A nuclide is an atom characterized by a definite atomic number and mass number. • The shorthand notation for a nuclide consists of its symbol with the atomic number, Z, as a subscript on the left and its mass number, A, as a superscript on the left. + 11p, 12 n, 11e- 11p, 12 n, 10e-

  14. What’s the element? 17 p  atomic number on periodic table for chorine What is the nuclide symbol for a nucleus that consists of 17 protons, 18 neutrons, and 17 electrons? How many protons, neutrons, and electrons are in the following nucleus A = #p + #n = 17 + 18 = 35 Note: #e- = #p; therefore, neutral species 35 p 45 n 36 e- A = #p + #n 35e- + one additional e- based on -1 charge = 36 e- 80 = 35 + #n #n = 80 – 35 = 45 n code for both:proton HW 9 HW 10 Note: #e- > #p; therefore, negatively charged species

  15. Atomic Theory of Matter • Isotopesare atoms whose nuclei have the same atomic number (Z) but different mass numbers (A); that is, the nuclei have the same number of protons but different numbers of neutrons thereby causing them to have different masses. • Chlorine, for example, exists as two isotopes: chlorine-35 and chlorine-37. Frac abund = 24.229% Mass = 36.97 amu Frac abund = 75.771% Mass = 34.97 amu 17p, 20 n, 17e- • The fractional abundanceis the fraction of a sample of atoms that is composed of a particular isotope. 17p, 18 n, 17e- (0.75771)(34.97 amu) + (0.24229) (36.97 amu) = 35.45 amu Note: The mixture of isotope masses make up the actual mass of the element given on periodic table. Cl has a mass of 35.45 amu which is based on the two isotopes of Cl-35 and Cl-37.

  16. Atomic Weights Calculate the atomic weight of boron, B, from the following data: ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE B-10 10.013 0.1978 (19.78%) B-11 11.009 0.8022 (80.22%) Note: fractional abundances must add to 1 (100%) B-10: 10.013 amu x 0.1978 = 1.980 B-11: 11.009 amu x 0.8022 = 8.831 10.811 = 10.811 amu ( = atomic wt.) Note: mass on periodic table matches 10.811 amu (weighted average of isotopes) code: amu HW 11

  17. Atomic Weights Dalton’s Relative Atomic Masses • Since Dalton could not weigh individual atoms, he devised experiments to measure their masses relative to the hydrogen atom. • Hydrogen was chosen as it was believed to be the lightest element. Daltons assigned hydrogen a mass of 1 (1 Dalton = mass of H). • For example, he found that carbon weighed 12 times more than hydrogen. He therefore assigned carbon a mass of 12 ( mass of carbon = 12 Daltons).

  18. Atomic Weights Dalton’s Relative Atomic Masses • Dalton’s atomic weight scale was eventually replaced in 1961, by the present carbon–12 mass scale. • One atomic mass unit (amu)is, therefore, a mass unit equal to exactly 1/12 the mass of a carbon–12 atom. • On this modern scale, the atomic weightof an element is the average atomic mass for the naturally occurring element, expressed in atomic mass units. Periodic table is based on atom mass with units of amu. Na - 23.1 amu mass of 1 atom of sodium

  19. The Periodic Table In 1869, Dmitri Mendeleev discovered that if the known elements were arranged in order of atomic mass (A), they could be placed in horizontal rows such that the elements in the vertical columns had similar properties. • periodic table - tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of the elements. • periodic law – states that certain sets of physical and chemical properties recur at regular intervals (periodically) when the elements are arranged according to increasing atomic number (Z). • Note: eventually changed from atomic mass to atomic number because of a couple of anomalies.

  20. Figure: A modern form of the periodic table. anomalies

  21. The Periodic Table Periods and Groups • A period consists of the elements in any one horizontal row of the periodic table. • A group consists of the elements in any one column of the periodic table (similar properties/structure). • The groups are usually numbered (North American uses roman numbers and A/B; IUPAC 1-18). • The eight “A” groups are called main group(or representative) elements.

  22. The Periodic Table • Periods and Groups • The “B” groups are called transitionelements. • The two rows of elements at the bottom of the table are called inner transition elements. • Elements in any one group have similar properties because their outer shells have the same number of valence electron (discuss in later sections).

  23. The Periodic Table • The elements in group IA (except H) - alkali metals • Periods and Groups • The elements in group IIA - alkaline earth metals, • The group VIIA elements - halogens • The group VIIIA elements – noble gases (monoatomic) • Diatomic elements – H2, N2, O2, F2, Cl2, Br2, I2 • Most species are solids at room temperature; H2, N2, O2, F2, Cl2, and noble gases are gases; Br2 and Hg are liquids.

  24. Metals, Nonmetals, and Metalloids – generally, left of staircaseare metals, touching staircase are metalloids, right of staircase are nonmetals. This is important for determining bond type, using proper terminology, and making decisions. Metallic character HW 12 nonmetals Metallic character metals code: table

  25. Chemical Formulas; Molecular and Ionic Substances The chemical formula of a substance is a notation using atomic symbols with subscripts to convey the relative proportions of atoms of the different elements in a substance. • aluminum oxide, Al2O3 2Al:3O ratio • sodium chloride, NaCl 1Na:1Cl ratio • calcium nitrate, Ca(NO3)2 1Ca:2NO3- ratio • or1Ca:2N:6O ratio

  26. Chemical Formulas; Molecular and Ionic Substances involves covalent bond – share electrons between atoms – typically nonmetal/nonmetal involves ionic bond – transfer electrons between atoms – attraction between charged particles – typically metal/nonmetal or polyatomic ions C : C Na+Cl- • Molecular substances • A molecule is a definite group of atoms that are chemically bonded together through sharing of electrons (covalent bonding, generally nonmetal-nonmetal including H). • A molecular substance is a substance that is composed of molecules, all of which are alike. • A molecular formula gives the exact number of atoms of elements in a molecule (i.e. C2H6O). • Structural formulasshow how the atoms are bonded to one another in a molecule. • i.e. ethanol (C2H6O) has a structural formula of CH3CH2OH

  27. Although many substances are molecular, others are composed of ions (charged particles) thathavetransferred electrons and have ionic bonding; occurs generally with metal-nonmetalinteractions. Ionic substances • An ion is an electrically charged particle obtained from an atom or chemically bonded group of atoms by adding or removing electrons. • Sodium chloride is a substance made up of ions. Na Cl + 1e- - Na+Cl-

  28. Chemical Formulas; Molecular and Ionic Substances Ionic substances • The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. • The formula unit of the substance is the group of atoms or ions explicitly symbolized by its formula. C : O Na+Cl-

  29. When an atom gains extra electrons, it becomes a negatively charged ion, called an anion (more electrons than protons). i.e, Cl- • An atom that loses electrons becomes a positively charged ion, called a cation (more protons than electrons). i.e., Na+ • An ionic compound is a compound composed of cations and anions. • NaCl • CaBr2 • Na2SO4 • CO2 Ionic substances Answer the following questions for species below: ionic or molecular substance; formula unit or molecule; ionic or covalent bonds involved? ionic substance; formula unit; ionic bond ionic substance; formula unit; ionic bonds ionic substance; formula unit; ionic and covalent bonds in SO42- molecular substance; molecule; covalent bonds

  30. Ions in Aqueous Solution Many (not all) ionic compounds (ionic bond/m-nm) dissociate into independent ions when dissolved in water NaCl (s)  Na+(aq) + Cl-(aq) Soluble ionic compounds dissociate 100% - referred to as strong electrolytes – breaks into charged particles until reaches saturation point. Soluble salt charges particles

  31. Ions in Aqueous Solution Most molecular (covalent bond/nm-nm) compounds dissolve but do not dissociate into ions, exception acids. C6H12O6 (s)  C6H12O6 (aq) These compounds are referred to as nonelectrolytes; no charged particles; soluble to saturation point but no ions formed. no charges particles; remains whole How would sodium sulfate dissolve based on bonding? Na2SO4 (s)  2Na+(aq) + SO42-(aq) ionic bond dissociates while covalent bonds in sulfate remain intact

  32. Chemical Substances; Formulas and Names Ionic compounds • Most ionic compounds contain metal and nonmetal atoms (as well as polyatomic ions); for example, NaCl. • You name an ionic compound by giving the name of the cation followed by the name of the anion with -ide. • Sodium chloride, NaCl Calcium Iodide, CaI2 • Potassium Bromide, KBr • We give the monatomic ion name for the cations and anions when naming compounds. A monatomic ion is an ion formed from a single atom.

  33. Most of the main group metals form cations with the charge equal to their roman group number. • The charge on a monatomic anion for a nonmetal equals the roman group number minus 8. • Most transition elements form more than one ion, each with a different charge (exceptions Cd2+, Zn2+, Ag+). • Other important elements with variable charge • Pb4+, Pb2+ Sn4+, Sn2+ As5+, As3+ Sb5+, Sb3+ How do we get the charge for ions? Rules for predicting charges on monatomic ions 0 1+ 4- 2+ 3+ 4+ 3- 2- 1- varies

  34. Monatomic cations are named after the element. For example, Al3+ is called the aluminum ion. • If there is more than one cation of an element (charge), a Roman numeral in parentheses denoting the charge on the ion is used. This often occurs with transition elements. • Na+ sodium ion Ca2+ calcium ion • Fe2+ iron (II) ion Fe3+ iron (III) ion • Older name: higher ox state (charge) – ic, / lower, -ous • Fe3+ferric ion Fe2+ferrous ion Cu2+cupric ion • Cu+ cuprous ion Hg2+mercuric ion Hg22+mercurousion • also done with Pb4+, Pb2+; Sn4+, Sn2+; As5+, As3+; Sb5+, Sb3+. • For the names of the monatomic anions, use the stem name of the element followed by the suffix – ide. For example bromine, the anion is called bromide ion, Br-. Rules for naming monatomic ions

  35. Based on the charge of the ions and balancing the overall charge on the compound by adjusting the number of ions, a formula is written. Note the sum of all the charges must equal zero, and you do not display the charges in the final formula. formula ions and charges NaCl Roman number tells charge of transition metal 1Na+= 1+ charge 1Cl- = 1- charge balanced The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. Sodium chloride Na+Cl- Iron (III) sulfate Fe3+ SO42- Chromium (III) oxide Cr3+O2- Calcium nitrate Ca2+ NO3- Sodium phosphate Na+ PO43- Strontium oxide Sr2+ O2- Fe2(SO4)3 2 3 2Fe3+ = 6+ charge 3SO42- = 6- charge balanced Cr2O3 Ca(NO3)2 Generally, you can crisscross the charge of one ion as the subscript on the second ion, reducing when possible. Na3PO4 HW 13 & 14 code for both: formula SrO

  36. Naming Ionic Binary Compounds To name a compound, you must know if it is a molecular or ionic compound so that you know which rules to follow. If you have a metal-nonmetal (or polyatomic ion), it is an ionic compound where you name the metal first then the nonmetal with changing the ending to –ide. If it is a transition metal, you must include the charge of the metal (Roman numbers). If you have nonmetal-nonmetal, it is a molecular compound which we haven’t discussed yet. sodium fluoride NaF - - lithium chloride MgO - MnBr2 - - cobalt (III) oxide - copper (II) chloride or cupric chloride LiCl magnesium oxide The charge on Mn must be 2+ to balance out the 2Br- charges. manganese (II) bromide The Roman number 3 tells us the charge on Co is 3+ which helps us determine the formula knowing that O is 2-. Co2O3 CuCl2

  37. Chemical Substances; Formulas and Names • A polyatomic ionis an ion consisting of two or more atoms chemically bonded together and carrying a net electric charge. We name the compounds the same way we just discussed except each polyatomic ion has a particular name. • Books typically have a table that lists common polyatomic ions. Most are oxo anions – consists of oxygen with another element (central element). Polyatomic ions NO3- nitrate SO42- sulfate NO2- nitrite SO32- sulfite Most groups have –ate, -iteendings and differ by #O. Mn, Br, Cl, I have per- -ate, -ate, -ite, hypo- -ite.

  38. O22- - Peroxide PO43- - Phosphate PO33- - Phosphite CO32- - Carbonate HCO3- - Bicarbonate or Hydrogen Carbonate N3- - azide NO3- - nitrate NO2- - nitrite C2H3O2-or CH3COO-- acetate Cr2O72- - dichromate CrO42- - chromate C2O42- - oxalate HSO4- - bisulfate or hydrogen sulfate H2PO4- - dihydrogen phosphate Ions You Should Know Polyatomic ions • NH4+ - Ammonium • OH- - Hydroxide • CN- - Cyanide • SO42- - Sulfate • SO32- - Sulfite • ClO4- - perchlorate • ClO3- - chlorate • ClO2- - chlorite • ClO- - hypochlorite • Hg22+ - mercury (I) or mecurous • S2O32- - thiosulfate • SCN- - thiocyanate • CNO- - cyanate • MnO4- - permanganate

  39. tin (II) sulfate or stannous sulfate Na2SO3 calcium hypochlorite SnSO4 sodium sulfite Ca(ClO)2 barium hydroxide potassium perchlorate Cr2(SO4)3 magnesium nitride Fe3(PO4)2 titanium (IV) nitrate Ba(OH)2 KClO4 chromium (III) sulfate Note: Not a polyatomic ion; monoatomic anion of N. Mg3N2 iron (II) phosphate or ferrous phosphate Ti(NO3)4

  40. Chemical Substances; Formulas and Names Molecular compounds • Binary compounds composed of two nonmetals are usually molecular and are named using a prefix system (name same as ionic except must indicate how many atoms are present using mono, di, tri, etc.). No charges (share electrons) involved with molecular compounds, but we typically put more metallic compound first. • Which way is the correct way to write the following formula based on putting the more metallic compound first? NF3 F3N

  41. Chemical Substances; Formulas and Names Binary molecular compounds • The name of the compound has the elements in the order given in the formula. • You name the first element using the exact element name. • Name the second element by writing the stem name of the element with the suffix “–ide.” • If there is more than one atom of any given element, you add a prefix (di, tri, tetra, penta, hexa, hepta, octa, etc.)

  42. N2O3 • SF4 • chlorine dioxide • sulfur hexafluoride • Cl2O7 • HCl (g) • Name this compound but think about bonding: • MgCl2 • Older names: water - H2O, ammonia – NH3, • hydrogen sulfide – H2S, nitric oxide – NO, hydrazine – N2H4 To name a compound, you must know if it is a molecular or ionic compound so that you know which rules to follow. If you have a metal-nonmetal (or polyatomic ion), it is an ionic compound where you name the metal first then the nonmetal with changing the ending to –ide. If you have nonmetal-nonmetal, it is a molecular compound which you do similarly as the ionic compound except that you must use prefixes to indicate the number of atoms. dinitrogen trioxide sulfur tetrafluoride • Binary molecular compounds ClO2 SF6 Drop the “a” on prefix if you encounter double vowel in name. dichlorineheptoxide Since this is a gas, we name using molecular rules; however, if acid we have other rules. hydrogen chloride magnesium chloride; ionic comp, no prefix

  43. Chemical Substances; Formulas and Names Acids • Acids are traditionally defined as compounds with a potential H+ as the cation. • Binary acids consist of a hydrogen ion and any single anion in aqueous solution. For example, HCl (aq) is hydrochloric acid. Binary acid: hydrostemicacid • An oxoacidis an acid containing hydrogen, oxygen, and another element. An example is HNO3, nitric acid. The oxoacids are a derivation of the oxoanions we discussed earlier.

  44. If you learn the oxoanions, you can easily adapt to naming the oxoacids: -ate  -ic and –ite  -ous oxoacids Anion prefix/suffixacid prefix/suffic per- -ate ion per- -ic acid -ate ion -ic acid -ite ion -ous acid hypo- -ite ion hypo- -ous acid NO3- nitrate ion HNO3 nitric acid NO2- nitrite ion HNO2 nitrous acid ClO4- perchlorate ion HClO4perchloric acid For some species there is a change in spelling in the name. SO42- sulfate ion H2SO4 sulfuric acid PO43- phosphateion H3PO4 phosphoric acid

  45. Chemical Substances; Formulas and Names • A hydrateis a compound that contains water molecules weakly bound in its crystals. Hydrates • Hydrates are named from the anhydrous (dry) compound, followed by the word “hydrate” with a prefix to indicate the number of water molecules per formula unit of the compound. • CuSO4. 5H2O • Magnesium sulfate heptahydrate copper(II)sulfate pentahydrate MgSO4.7H2O HW 15 - 18 code for all: names

  46. Chemical Substances; Formulas and Names Naming simple compounds • Chemical compounds are classified as organic or inorganic. • Organic compounds are compounds that contain carbon combined with other elements, such as hydrogen, oxygen, and nitrogen. • Inorganic compounds are compounds composed of elements other than carbon.

  47. Chemical Formulas; Molecular and Ionic Substances Organic compounds • An important class of molecular substances that contain carbon is the organic compounds. • Organic compounds make up the majority of all known compounds. • The simplest organic compounds are hydrocarbons - compounds containing only hydrogen and carbon. • Common examples include methane, CH4, ethane, C2H6, and propane, C3H8.

  48. Classifying CompoundsOrganic vs. Inorganic • in the 18th century, compounds from living things were called organic; compounds from the nonliving environment were called inorganic • organic compounds easily decomposed and could not be made in 18th century lab • inorganic compounds very difficult to decompose, but able to be synthesized

  49. Modern Classifying CompoundsOrganic vs. Inorganic • today we commonly make organic compounds in the lab and find them all around us • organic compounds are mainly made of C and H, sometimes with O, N, P, S, and trace amounts of other elements • the main element that is the focus of organic chemistry is carbon

  50. Carbon Bonding • carbon atoms bond almost exclusively covalently • compounds with ionic bonding C are generally inorganic • when C bonds, it forms 4 covalent bonds • 4 single, 1 double + 2 singles, 2 double, or 1 triple + 1 single • carbon is unique in that it can form limitless chains of C atoms, both straight and branched, and rings of C atoms

More Related