1 / 50

Chapter Three

Fundamentals of General, Organic and Biological Chemistry 6th Edition. Chapter Three. Atoms and the Periodic Table. Outline. 3.1 Atomic Theory 3.2 Elements and Atomic Number 3.3 Isotopes and Atomic Weight 3.4 The Periodic Table 3.5 Some Characteristics of Different Groups

patsy
Télécharger la présentation

Chapter Three

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. Fundamentals of General, Organic and Biological Chemistry 6th Edition Chapter Three Atoms and the Periodic Table

  2. Outline • 3.1 Atomic Theory • 3.2 Elements and Atomic Number • 3.3 Isotopes and Atomic Weight • 3.4 The Periodic Table • 3.5 Some Characteristics of Different Groups • 3.6 Electronic Structure of Atoms • 3.7 Electron Configurations • 3.8 Electron Configurations and the Periodic Table • 3.9Electron-Dot Symbols Chapter Three

  3. Goals • 1. What is the modern theory of atomic structure?Be able to explain the major assumptions of atomic theory. • 2. How do atoms of different elements differ?Be able to explain the composition of different atoms according to the number of protons, neutrons, and electrons they contain. • 3. What are isotopes, and what is atomic weight?Be able to explain what isotopes are and how they affect an element’s atomic weight. Chapter Three

  4. Goals Contd. • 4. How is the periodic table arranged?Be able to describe how elements are arranged in the periodic table, name the subdivisions of the periodic table, and relate the position of an element in the periodic table to its electronic structure. • 5. How are electrons arranged in atoms?Be able to explain how electrons are distributed in shells and subshells around the nucleus of an atom, how valence electrons can be represented as electron-dot symbols, and how the electron configurations can help explain the chemical properties of the elements. Chapter Three

  5. 3.1 Atomic Theory • Chemistry is founded on four fundamental assumptions about atoms and matter, which together make up modern Atomic Theory: • 1. All matter is composed of atoms. • 2. The atoms of a given element differ from the atoms of all other elements. • 3. Chemical compounds consist of atoms combined in specific ratios. • 4. Chemical reactions change only the way the atoms are combined in compounds; the atoms themselves are unchanged. Chapter Three

  6. Atoms are composed of tiny subatomic particles called protons, neutrons, and electrons. • The masses of atoms and their constituent subatomic particles are very small when measured in grams. Atomic masses are expressed on a relative mass scale. One atom is assigned a mass, and all others are measured relative to it. • The basis for the relative atomic mass scale is an atom of carbon that contains 6 protons and 6 neutrons. This carbon atom is assigned a mass of exactly 12 atomic mass units (amu). Chapter Three

  7. Both protons and neutrons have a mass close to one amu. • Electrons are 1836 times lighter than protons and neutrons. • Protons and electrons have electrical charges that are equal in magnitude but opposite in sign. Chapter Three

  8. The protons and neutrons are packed closely together in a dense core called the nucleus. Surrounding the nucleus, the electrons move about rapidly through a large volume of space. • The relative size of a nucleus in an atom is the same as that of a pea in the middle of this stadium. Chapter Three

  9. Diameter of a nucleus is only about 10-15 m. • Diameter of an atom is only about 10-10 m. Chapter Three

  10. The structure of the atom is determined by interplay of different forces. • Opposite electrical charges attract each other, like charges repel each other. • Protons and neutrons in the nucleus are held together by the nuclear strong force. Chapter Three

  11. 3.2 Element and Atomic Number • Atomic Number(Z): The number of protons in each atom of an element. All atoms of a particular element have the same number of protons in the nucleus. • Atoms are neutral overall and have no net charge because the number of positively charged protons and the number of negatively charged electrons are the same in each atom. • Mass Number(A):The total number of protons and neutrons in an atom. Chapter Three

  12. Atomic Notation Isotope Mass number (A) • A = (# protons) + (#neutrons) • A = Z + N • For charge neutrality, number of electrons & protons must be equal Atomic Symbols • Symbolized by

  13. Learning Check An atom of has ___ protons, ___ neutrons, and ___ electrons. • 82, 206, 124 • 124, 206, 124 • 124, 124, 124 • 82, 124, 82 • 82, 124, 124

  14. 3.3 Isotopes and Atomic Weight • Isotopes: Atoms with identical atomic numbers (Z) but different mass numbers (A) are called isotopes. Protium, deuterium, and tritium are three isotopes of the element hydrogen. • H, most abundant hydrogen isotope has one proton and no neutrons (Z=1, A=1) • D, this heavy hydrogen isotope has one proton and one neutron (Z=1, A=2), • T, this radioactive hydrogen isotope has one proton and two neutrons (Z=1, A=3). Chapter Three

  15. Isotopes of Hydrogen Chapter Three

  16. A specific isotope is represented by showing its mass number (A) as a superscript and its atomic number (Z) as a subscript in front of the atomic symbol. For • example, the symbol for tritium is: Chapter Three

  17. Atomic Weight: The weighted average mass of an element’s atoms in a large sample that includes all the naturally occurring isotopes of that atom. • To calculate the atomic weight of an element, the individual mass and the percent abundance of each naturally occurring isotope must be known. • Atomic weight= (isotope abundance)*(isotope mass) • The Greek symbol, , indicates the summation of terms over all naturally occurring isotopes. Chapter Three

  18. Calculating Atomic Mass • Generally, elements are mixtures of isotopes Ex. Hydrogen Isotope Mass %Abundance 1H 1.007825 amu 99.985 2H 2.0140 amu 0.015 How do we define Atomic Mass? • Average of masses of all stable isotopes of given element How do we calculate Average Atomic Mass? • Weighted average. • Use Isotopic Abundances & isotopic masses

  19. Learning Check Naturally occurring magnesium is a mixture of 3 isotopes; 78.99% of the atoms are 24Mg (atomic mass, 23.9850 amuu), 10.00% of 25Mg (atomic mass, 24.9858 amu), and 11.01% of 26Mg (atomic mass, 25.9826 amu). From these data calculate the average atomic mass of magnesium. 0.7899 * 23.9850 amu = 18.946 amu 24Mg 0.1000 * 24.9858 amu =2.4986 amu 25Mg 0.1101 * 25.9826 amu =2.8607amu26Mg Total mass of average atom = 24.3053 amu rounds up to 24.31 amu

  20. Your Turn! A naturally occurring element consists of two isotopes. The data on the isotopes: isotope #1 68.5257 amu 60.226% isotope #2 70.9429 amu 39.774% Calculate the average atomic mass of this element. • 70.943 u • 69.487 u • 69.526 u • 69.981u • 69.734 u 0.60226 * 68.5257 amu = 41.270 amu 0.39774 * 70.9429 amu = 28.217 amu 69.487 amu

  21. 3.4 The Periodic Table • Beginning at the upper left corner of the periodic table, elements are arranged by increasing atomic number into seven horizontal rows, called periods, and 18 vertical columns, called groups. • The elements in a given group have similar chemical properties. Lithium, sodium, potassium and other elements in group 1A (or 1) have similar properties. Similarly, chlorine, bromine, iodine, and other elements in group 7A (or 17) behave similarly. Chapter Three

  22. The Periodic Table Chapter Three

  23. The table has 114 boxes, each of which tells the symbol, atomic number, and atomic weight of an element. Chapter Three

  24. All seven periods do not contain the same number of elements. • The first period contains only 2 elements. • The second and third periods each contains 8 elements. • The fourth and fifth periods each contain 18 elements. • The sixth period contains 32 elements. • The seventh period, still incomplete Chapter Three

  25. The groups on the periodic table are divided into three main categories. • Main Groups: The two groups on the far left (1-2) and the six on the far right (13-18) are the main groups. • Transition Metal Groups: Elements in the groups numbered 3 through 12. • Inner Transition Metal Groups:The 14 groups shown at the bottom of the table that are not numbered containing the Lanthanides and the Actinides. Chapter Three

  26. 3.5 Characteristics of Different Groups • Group 1A or 1 Alkali metals: • Li, Na, K, Rb, Cs, and Fr • Shiny, soft, and low melting point metals • All react rapidly with water to form flammable H2 gas and alkaline or basic solutions Chapter Three

  27. Group 2A or 2 Alkaline earth metals: • Be, Mg, Ca, Sr, Ba, and Ra • Lustrous, silvery metals • React with O2 • They are less reactive to water than the alkali metals Chapter Three

  28. Group 7A or 17 Halogens: • F, Cl, Br, I, and At • Colorful and corrosive nonmetals • All are found in nature in combination with other elements, such as with sodium in sodium chloride (NaCl) Chapter Three

  29. Group 8A or 18 Noble gases: • He, Ne, Ar, Kr, Xe, and Rn • Colorless gases • Very low chemical reactivity Chapter Three

  30. 3. 6 Electronic Structure of Atoms Quantum mechanical model of atomic structure: • The electrons in an atom are grouped around the nucleus into shells, roughly like the layers in an onion. • The farther a shell is from the nucleus, the larger it is, the more electrons it can hold, and the higher the energies of those electrons. • The smallest shell closest to the nucleus is labeled shell 1,the next one is shell 2, and so on. Chapter Three

  31. Within the shells, electrons are further grouped into subshells of four different types, identified as s, p, d, and f in order of increasing energy. • A shell has a number of subshells equal to its shell number. • The first shell has only an s subshell; the second shell has an s and a p subshell; the third shell has an s, a p, and a d subshell, and so on. Chapter Three

  32. Within each subshell, electrons are further grouped into orbitals, regions of space within an atom where the specific electrons are more likely to be found. • The number of orbitals within a subshell increases as the odd numbers. • An s subshell has 1 orbital, a p has 3, a d has 5 and so on. Chapter Three

  33. Different orbitals have different shapes. Orbitals in s subshells are spherical (a), while orbitals in p subshells are roughly dumbbell shaped (b). Chapter Three

  34. Any orbital can hold a maximum of 2 electrons. • The first shell has one 1s orbital and holds 2 electrons. • The second shell can hold 8 electrons, 2 in a 2s orbital and 6 in three 2p orbitals. • The third shell can hold 18 electrons, 2 in a 3s orbital, 6 in three 3p orbitals, and 10 in five 3d orbitals, and so on. Chapter Three

  35. The overall electron distribution within an atom is summarized in table 3.2 below. Chapter Three

  36. 3.7 Electron Configurations • Electron Configuration: The exact arrangement of electrons in atom’s shells and subshells. Rules to predict electron configuration: • Electrons occupy the lowest energy orbitals available first. • Each orbital can hold only two electrons, which must be of opposite spin. • If two or more orbitals have the same energy, each orbital gets one electron before any orbital gets two. Chapter Three

  37. Order of orbital energy levels: • Electrons fill orbitals from the lowest-energy orbitals upward. • Lower numbered shells fill before higher numbered shells at first. • Some overlap in energy levels occurs starting with shell 3 and 4. Chapter Three

  38. Below is a mnemonic device for remembering the order of the energy levels in an atom. Chapter Three

  39. Order of orbitals (filling) in multi-electron atom 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

  40. Electron configurations are described by writing the shell number and subshell letter in order of increasing energy. The number of electrons actually occupying each subshell is indicated by a superscript. • A graphic representation can be made by indicating each orbital as a line and each electron as an arrow. The head of the arrow indicates the electron spin. • A shorthand using noble gas configurations is very useful for large atoms. Chapter Three

  41. These are the electron configurations for B - N Chapter Three

  42. These are the electron configurations for O - Ne Chapter Three

  43. What is the electron configuration of Mg? What is electron configuration of Cl? Mg 12 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s2 2 + 2 + 6 + 2 = 12 electrons Abbreviated as [Ne]3s2 [Ne] 1s22s22p6 Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electrons Last electron added to 3p orbital

  44. 3.8 Electron Configuration and the Periodic table • Valence Shell: Outermost, highest energy shell of an atom. • Valence electrons: An electron in an outermost shell of an atom. These electrons are loosely held, they are most important in determining an element’s properties. Chapter Three

  45. The periodic table can be divided into four regions or blocks, of elements according to the subshells that are last to fill, s, p, d, or f. • Beginning at the top left corner of the periodic table, the first row contains only two elements H and He. The 1s subshell is being filled here. • The second row begins with two s-block elements (Li and Be) and continues with six p-block elements (B through Ne), so electrons fill the next available s orbital (2s) and then the first available p orbitals (2p). Chapter Three

  46. The third row is similar to the second row, so the 3s and 3p orbitals are filled next. • The fourth row again starts with two s-block elements (K and Ca) but is then followed by ten d-block elements (Sc through Zn) and six p-block elements (Ga through Kr). Thus, the order of orbital filling is 4s followed by the first available d orbitals (3d) followed by 4p. • Continuing through successive rows of the periodic table provides a visual method to recall the entire filling order. Chapter Three

  47. Visual method to recall the order of orbital filling. Chapter Three

  48. Analysis WORKED EXAMPLE 3.7Electron Configurations: Orbital-Filling Diagrams Write the electron configuration of phosphorus, using up and down arrows to show how the electrons in each orbital are paired. Phosphorus has 15 electrons, which occupy orbitals according to the order shown in Figure 3.6. • The first 2 are paired and fill the first shell (1s2). • The next 8 fill the second shell (2s2 2p6). All electrons are paired. • The remaining 5 electrons enter the third shell, where 2 fill the 3s orbital (3s2) and 3 occupy the 3p subshell, one in each of the three p orbitals. Solution

  49. Analysis (a) Na (sodium) is located in the third row, and in the first column of the s-block. Therefore, all orbitals up to the 3s are completely filled, and there is one electron in the 3s orbital. (b) Cl (chlorine) is located in the third row, and in the fifth column of the p-block. (c) Zr (zirconium) is located in the fifth row, and in the second column of the d-block. All orbitals up to the 4d are completely filled, and there are 2 electrons in the 4d orbitals. Note that the 4d orbitals are filled after the 5s orbitals in both Figures 3.6 and 3.7. WORKED EXAMPLE 3.8Electron Configurations: Valence Electrons Write the electron configuration for the following elements, using both the complete and the shorthand notations. Indicate which electrons are the valence electrons. (a) Na          (b) Cl          (c) Zr Locate the row and the block in which each of the elements is found in Figure 3.7. The location can be used to determine the complete electron configuration and to identify the valence electrons. Solution

  50. Chapter Three 3.9 Electron-Dot Symbols • Electron-dot symbol: An atomic symbol with dots placed around it to indicate the number of valence electrons.

More Related