Presentation on theme: "1 Electronic Structure of Atoms & Periodic Table Chapter 4 CHEMISTRY - DMCU 1233 Fakulti Kejuruteraan Mekanikal, UTeM Lecturer: IMRAN SYAKIR BIN MOHAMAD."— Presentation transcript:
1 Electronic Structure of Atoms & Periodic Table Chapter 4 CHEMISTRY - DMCU 1233 Fakulti Kejuruteraan Mekanikal, UTeM Lecturer: IMRAN SYAKIR BIN MOHAMAD MOHD HAIZAL BIN MOHD HUSIN NONA MERRY MERPATI MITAN
2 4.1 HISTORY OF ATOMIC MODEL ContributorModelExplanation John Dalton (1805) Billiard Ball Model Daltons atomic model was represented as a small united ball similar to a very tiny ball. J. J Thomson (1897) Plum Pudding Model Thomson discovered the electron, a negatively charged particle. The atom was described as a sphere of positive charge with electrons embedded in it. Ernest Rutherford (1911) Solar System Model Rutherford discovered the proton, a positively charged particle in an atom. The proton and most of the mass of the atom were concentrated in the central region called the nucleus. The electrons moved in the spherical space outside the nucleus. Neils Bohr (1913) Bohr Model According to Bohr, the electrons in an atom were not randomly distributed around the atomic nucleus, but moved around the nucleus in fixed orbits (shell). Each orbit formed a circle and had a fixed distance from the nucleus.
3 4.1 HISTORY OF ATOMIC MODEL Billiard Ball Model Plum Pudding Model Solar System Model Bohr Model
4 4.2 Quantum Numbers The Bohr model was 1-D model that used one quantum number to describe the distribution of electrons in the atom that representative of the size of the orbit, which was described by the principal quantum number (n). Meanwhile Schrödinger's model allowed the electron to occupy in 3-D space to describe the orbitals in which electrons can be found. Each electron in an atom is described by four different quantum numbers. The first three quantum number from Schrödinger's wave equations are the principal (n), angular (l), and magnetic (m l ) quantum numbers describe the size, shape, and orientation in space of the orbitals on an atom. The fourth quantum number spin (m s ) specifies how many electrons can occupy that orbital. A quantum number describes the energies of electrons in atoms
5 4.2 Quantum Numbers Principal quantum number – ( n ) Angular momentum quantum number – ( l ) Magnetic quantum number – ( m l ) Spin quantum number – ( m s )
6 Quantum Numbers ( n, l, m l, m s ) Principal quantum number n n = 1, 2, 3, 4, …. n =1 n=2 n=3 Specifies the energy of an electron and the size of the orbital (the distance from the nucleus of the peak in a radial). All orbitals that have the same value of n are said to be in the same shell (level) The total number of orbitals for a given n value is n 2
7 Angular momentum quantum number l for a given value of n, l = 0, 1, 2, 3, … n-1 n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 Shape of the volume of space that the e - occupies l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital Quantum Numbers (n, l, m l, m s ) Specifies the shape of an orbital with a particular principal quantum number. The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels).
8 l = 0 (s orbitals) l = 1 (p orbitals)
9 l = 2 (d orbitals)
10 l = 3 (f orbitals)
11 Magnetic quantum number m l for a given value of lm l = -l, …., 0, …. +l orientation of the orbital in space for l = 0 (s orbital) m l = 0 if l = 1 (p orbital), m l = -1, 0, or +1 if l = 2 (d orbital), m l = -2, -1, 0, +1, or +2 Quantum Numbers ( n, l, m l, m s ) Specifies the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the subshell into individual orbitals which hold the electrons. There are 2l+1 orbitals in each subshell.
12 m l = -1m l = 0m l = 1 m l = -2m l = -1m l = 0m l = 1m l = 2
13 ORIENTATION OF THE ORBITAL IN SPACE
14 Spin quantum number m s m s = +½ or -½ Quantum Numbers ( n, l, m l, m s ) m s = -½m s = +½ Experimental arrangement for demo the spinning motion of electrons Q & A Specifies the orientation of the spin axis of an electron. An electron can spin in only one of two directions (sometimes called up and down).
15 Quantum number SymbolMeaning Range of values Value examples principalnshell1 nn = 1, 2, 3, … angular momentum subshell (s orbital is listed as 0, p orbital as 1 ) 0 n 1 for n = 3: = 0, 1, 2 (s, p, d) magneticm Orbital (orientation of the subshell's shape) m for = 2: m = 2, 1, 0, 1, 2 spinmsms spin of the electron (½ = "spin down", +½ = "spin up") s m s s for an electron s = ½, so m s = ½, +½
16 nlmlml Number of orbitals Orbital Name Number of electrons 10011s s2 1-1, 0, +132p s2 1-1, 0, +133p6 2-2, -1, 0, +1, +253d s2 1-1, 0, +134p6 2-2, -1, 0, +1, +254d10 3-3, -2, -1, 0, +1, +2, +374f14 Relation between quantum number, atomic orbital and number of an electron
17 Existence (and energy) of electron in atom is described by its unique Quantum Numbers Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945) No two electrons in the same atom can have identical values for all four of their quantum numbers Quantum Numbers ( n, l, m l, m s ) Two electrons in the same orbital must have opposite spins. Because an electron spins, it creates a magnetic field, which can be oriented in one of two directions. For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired
18 Paramagnetic unpaired electrons 2p Diamagnetic all electrons paired 2p The substances are not attracted to magnets and are said to be diamagnetic. Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic.
19 Quantum Numbers (n, l, m l, m s ) Shell – electrons with the same value of n Subshell – electrons with the same values of n and l Orbital – electrons with the same values of n, l, and m l How many electrons can an orbital hold?
20 How many 2p orbitals are there in an atom? How many electrons can be placed in the 3d subshell? Q & A
21 Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom. 1s 1 principal quantum number n angular momentum quantum number l number of electrons in the orbital or subshell Orbital diagram is shows the spin of the electron H 1s Electron configuration
22 Order of orbitals (filling) in multi-electron atom 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f Fill up electrons in lowest energy orbitals Aufbau principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states.
23 Outermost subshell being filled with electrons The order in which the electrons are filled in can be read from the periodic table in the following fashion
24 H 1 electronH 1s 1 Li 3 electronsLi 1s 2 2s 1 B 5 electronsB 1s 2 2s 2 2p 1 C 6 electronsC 1s 2 2s 2 2p 2
25 The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hunds rule). N 7 electrons N 1s 2 2s 2 2p 3 O 8 electronsO 1s 2 2s 2 2p 4 F 9 electronsF 1s 2 2s 2 2p 5 Ne 10 electronsNe 1s 2 2s 2 2p 6
26 What is the electron configuration of Mg? What are the possible quantum numbers for the last (outermost) electron in Cl?
27 Electron Configurations of Cations and Anions Na [Ne]3s 1 Na + [Ne] Ca [Ar]4s 2 Ca 2+ [Ar] Al [Ne]3s 2 3p 1 Al 3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s 1 H - 1s 2 or [He] F 1s 2 2s 2 2p 5 F - 1s 2 2s 2 2p 6 or [Ne] O 1s 2 2s 2 2p 4 O 2- 1s 2 2s 2 2p 6 or [Ne] N 1s 2 2s 2 2p 3 N 3- 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble- gas outer electron configuration. Of Representative Elements
Cations and Anions Of Representative Elements
29 Na + : [Ne]Al 3+ : [Ne] F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne]N 3- : 1s 2 2s 2 2p 6 or [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ?
30 Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5
31 Ion charges
32 Name the orbital described by the following quantum numbers : a.n = 3, l = 0 b.n = 3, l = 1 c.n = 3, l = 2 d.n = 5, l = 0 Q & A session
33 Give the n and l values for the following orbital a. 1s b. 3s c. 2p d. 4d e. 5f What and the possible m l values for the following types of orbital? a. s b. p c. d d. f Q & A session
34 How many possible orbital are there for n = a. 4 b. 10 How many electrons can inhabit all of the n = 4 orbital? Place the following orbital in order of increasing energy: 1s, 3s, 4s, 6s, 3d, 4f, 3p, 7s, 5d, 5p Q & A session
35 Write electron configurations for the following atoms: a. H b. Li + c. N d. F - e. Ca Q & A session
36 Draw an orbital diagrams for atoms with the following electron configurations: 1s 2 2s 2 2p 6 3s 2 3p 3 Q & A session
38 When the Elements Were Discovered
HISTORY OF THE PERIODIC TABLE 1.Antoine Lavoisier (1743–1794) Classify elements into four groups including light and heat, into metals and non-metals.
HISTORY OF THE PERIODIC TABLE 2.Johann Dobereiner (1780–1849) The first significant groupings of elements by place certain elements in groups of three known as The Law of Triads. Founded that strontium had about the average properties of calcium and barium, and grouped these three together. Several more triad groups, including the halogen triad of chlorine, bromine, and iodine, and the alkali metal triad of lithium, sodium, and potassium. However, due to the inaccuracy of many measurements, including atomic weight, the relationship between large element groups could not be exacted
HISTORY OF THE PERIODIC TABLE 3.John Newlands (1837–1898) arranged known elements horizontally in the ascending order of their atomic masses Each row consisted of seven elements. Founded that elements with similar properties repeated at every eighth element. This arrangement was known as the Law of Octaves However, this law was only obeyed by the first 17 elements. There were no positions allocated for elements yet to be discovered
HISTORY OF THE PERIODIC TABLE 4.Lothar Meyer (1830–1895) Plotted a graph of atomic volume against atomic mass for all known elements. Founded that elements with the same chemical properties occupied the same relative positions on the curve. Showed that the properties of the elements were in a periodic pattern with their atomic masses. Proved that the properties of the elements recur periodically.
HISTORY OF THE PERIODIC TABLE 5.Dmitri Mendeleev (1834–1907) Showed that the properties of elements changed periodically with their atomic mass. Arranged the elements in the order of increasing atomic mass and grouped them according to similar chemical properties. Able to predict the properties of undiscovered elements and left gap for these elements For examples correctly predicted the properties of the elements gallium, scandium and germanium which were only discovered later Mendeleevs table was used as a blueprint for the modern periodic table Mendeleevs periodic table
HISTORY OF THE PERIODIC TABLE 6.Henry J. G. Moseley (1887–1915) Based on the x-ray spectrum of elements studies, he concluded that the proton numbers should be used as a basis for the periodic change of chemical properties instead of the atomic mass. Rearranged the elements in the ascending order of their proton numbers Similar to Mendeleev, Moseley left gaps for elements yet to be discovered. produced a periodic table which was almost the same as Mendeleevs periodic table. Due to Moseleys work, the periodic table was successfully developed and being used today. The modern periodic table is based on the arrangement of elements in the ascending order of their proton numbers.
HISTORY OF THE PERIODIC TABLE 7.Glenn Seaborg discovered that the transuranium elements that have atomic numbers from 94 to 102, resulting in the redesign of the periodic table Technically, both the lanthanide and actinide series of elements are to be placed between the alkaline earth metal and the transition metal. However, by doing this, the periodic table would be too wide. Thus, the lanthanide and actinide series of elements were placed under the rest of the periodic table. Dr Seaborg and his colleagues were also responsible for identifying more than 100 isotopes of elements.
MODERN PERIODIC TABLE The periodic table is a systematic classification of elements whereby elements with the same chemical properties are placed in the same group. The elements in the periodic table are arranged in rows called the periods and columns which are known as the groups
MODERN PERIODIC TABLE Groups There are 18 groups of elements in the periodic table. Some of these groups have special names: (a) Group 1 elements are called alkali metals. (b) Group 2 elements are called alkaline earth metals. (c) Group 3 to Group 12 elements are known as transition elements. (d) Group 17 elements are called halogens. (e) Group 18 elements are called noble gases. Each member of a group shows similar chemical properties although their physical properties such as density, melting point and colour show a gradual change when descending the group.
MODERN PERIODIC TABLE Periods There are seven rows from period 1 to period 7. The elements are arranged horizontally in the ascending order of their proton numbers in the periodic table. The position of the period of an element in the periodic table is determined by the number of shells occupied with electrons in the atom. Period 1 has 2 elements only H and He, Periods 2 and 3 have 8 elements each. Periods 4 and 5, they have 18 elements each and they are called the long periods. Period 6 has 32 elements whereas the elements with proton number 58 to 71 are separated and are grouped below the periodic table known as the Lanthanide Series. Period 7 has 32 elements, the elements with proton number 90 to 103 are grouped below the periodic table known as the Actinide Series.
49 Periodic Classification Classification as metals and non-metals (a) Metals – a good conductor of heat and electricity. (b) Non-metals - a poor conductor of heat and electricity. (c) Metalloids – a intermediate between metal and non-metal properties
50 ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d1d1 d5d5 d 10 4f 5f Periodic Classification Ground State Electron Configurations of the Elements The similarity of the outer electron configuration (same type of valence electrons) makes the elements in the same group resemble one another in chemical behavior.
51 Periodic Classification Classification based on subshell filled with electron
52 Periodic Classification Representative elements (incompletely filled s and p subshells) Transition metals (incompletely filled d subshells) Noble gases (completely filled p subshells) Actinides (incompletely filled 5f subshells) Lanthanides (incompletely filled 4f subshells) Zn, Cd, Hg (neither representative element nor transition metals)
53 Periodic Trends Periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of the periodic table of elements Elemental characteristics a.Atomic radius b.Ionization energy c.Electron affinity d.Electronegativity e.Metallic properties f.Non-metallic properties
54 Periodic Trends Atomic radius The atomic radius is the distance between an atom's nucleus and its valence electrons in an atom. The atomic radius tends to decrease as one progresses across a period from left to right because the effective nuclear charge increases, thereby attracting the orbiting electrons and reducing the radius. The atomic radius usually increases while going down a group due to the addition of a new energy level (shell).
55 Periodic Trends Ionic Radius The ionic radius is different from the atomic radius of an element. Positive ions are smaller than their uncharged atoms. Negative ions are larger than their atoms.
56 Periodic Trends Ionization energy The ionization potential is the minimum amount of energy required to remove one electron from each atom. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius).
57 Periodic Trends Electron affinity The electron affinity described as the energy gained by an atom when an electron is added. Electron affinities becoming increasingly from left to right Electron affinities change little moving down a group, however they do generally become slightly more positive (less attractive toward electrons)
58 Periodic Trends Electronegativity Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. Electronegativity generally increases moving across a period and decreases moving down a group. Electronegativity plays a very large role in the processes of Chemical Bonding.
59 Periodic Trends Metallic properties Metallic property decreases across a period with increase in number of valence electrons as well as a decrease in atomic radius, and it increases down the group with increase in number of shells and atomic radius. Non-metallic properties Non-metallic property increases across a period and decreases down the group due to the same reason