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Elements & Compounds
Atoms (Elements)
Molecules (Compounds)
Cells
Elements & Compounds
Nucleus
Electrons Cloud of negative charge (2 electrons)
Fig. 2.5: Simplified model of a Helium (He) Atom
He 4.002602
2
Helium Mass Number (~atomic mass)
= number of Neutrons + Protons = 4 for Helium
Atomic number = number of protons
Atomic Mass
The Periodic Table of the Elements
Properties of Elements Depends Upon Atomic Structure
Atoms of different elements differ in their number of subatomic particles:
H 1 1 C 6
12 N 14 7
O 16 8 Na 11
23
C 6 12 C 6
13 Stable isotopes
C 6 14
Unstable (radioactive)
P 15 31
First shell
Second shell
Third shell
Hydrogen 1H
Lithium 3Li
Sodium 11Na
Beryllium 4Be
Magnesium 12Mg
Boron 5B
Aluminum 13Al
Carbon 6C
Silicon 14Si
Nitrogen 7N
Phosphorus 15P
Oxygen 8O
Sulfur 16S
Fluorine 9F
Chlorine 17Cl
Neon 10Ne
Argon 18Ar
Helium 2He
2 He
4.00 Mass number
Atomic number
Element symbol Electron distribution diagram
Fig. 2.9: Electrons are distributed in shells of orbitals. Each orbital contains a maximum of two electrons.
Chemical behavior of an atom depends mostly on number of electrons in outermost shell called Valence Electrons
Fig. 2.8: Energy Levels of Electrons / Electron Shells
Electrons in an atom vary in the amount of potential energy they possess (fixed, discrete amounts)
Different discrete energy levels correlate with average distance of electron from nucleus (electron shells)
Lower energy
Higher energy
(a) A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons
We Are Carbon-Based Lifeforms
Trace elements (< 0.01%): B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn
~4 %: Ca, P, K, S, Na, Cl, Mg
Composition of other organisms on Earth is similar
~96 % of the Human Body is: Oxygen (O) Carbon (C)
Nitrogen (N) Hydrogen (H)
Fig. 2.10: Electrons are arranged in orbitals that have characteristic shapes and energies
Neon, with two filled Shells (10 electrons)
First shell
Second shell
First shell Second shell
1s orbital 2s orbital Three 2p orbitals
(a) Electron distribution diagram
(b) Separate electron orbitals
(c) Superimposed electron orbitals
1s, 2s, and 2p orbitals
x y
z
Fig. 2.11: Chemical Bonds Link Atoms Together 1H1
1H1
H2
Bonding properties of the most common elements in biological molecules
H: Atomic No. 1 1 electron: 1s1
1st shell needs 1 electron Forms 1 covalent bond
C: Atomic No. 6 6 electrons: 1s2 2s2 2px1 2py
1 2pz0
2nd shell needs 4 electrons Forms 4 covalent bonds
N: Atomic No. 7 7 electrons: 1s2 2s2 2px1 2py
1 2pz1
2nd shell needs 3 electrons Forms 3 covalent bonds
O: Atomic No. 8 8 electrons: 1s2 2s2 2px2 2py
1 2pz1
2nd shell needs 2 electrons Forms 2 covalent bonds
Fig. 2.17a: Hybrid Atomic Orbitals
s orbital z
x
y
Three p orbitals
Hybridization of orbitals Tetrahedron
(a)
Space-Filling Model
Ball-and-Stick Model
Hybrid-Orbital Model (with ball-and-stick
model superimposed) Unbonded Electron pair
Water (H2O)
Methane (CH4)
(b) Molecular-shape models
Fig. 2.17b: Molecular Shape Models
LE 2-11 Fig. 2.12: Covalent Bonding in Four Molecules
(a) Hydrogen (H2)
(b) Oxygen (O2)
(c) Water (H2O)
Name and Molecular Formula
Electron Distribution
Diagram
Lewis Dot Structure and
Structural Formula
Space- Filling Model
(d) Methane (CH4)
• The two C atoms in Ethene form a second or double bond between themselves using their third unhybridized 2p orbitals.
• The second bond is called a π bond and prevents the C atoms from rotating around the bonds connecting them. • Note: the π bond does not place electron density between the atoms.
π bond
Element Symbol Electronegativity
Oxygen O 3.5
Chlorine Cl 3.1
Nitrogen N 3.0
Carbon C 2.5
Phosphorous P 2.1
Hydrogen H 2.1
Sodium Na 0.9
Potassium K 0.9
Chemical Bonds: Linking Atoms Together
High
High
O H H
δ-
δ+ δ+
Partial negative charge on oxygen
δ-
δ+ Partial positive charge on each hydrogen δ+ H2O
H H
O
Fig. 2.13: Polar Covalent Bonds Are Formed Between Atoms With Unequal Electronegativity
δ-
δ-
δ+ δ+ 104.5°
Fig. 2.17b: sp3 Hybrid Orbitals in Water, H2O
Fig. 3-2
Hydrogen bond
δ – H
δ + H
O —
—
δ + δ +
δ +
δ –
δ –
δ – • Hydrogen Bonds are attractions between the partial negative charge on the Oxygen atoms and the partial positive charges on the Hydrogen Atoms.
• This gives water special properties that we will discuss next time.
LE 2-15
δ–
Water (H2O)
Ammonia (NH3)
δ+
δ+
δ–
δ+
δ+
δ+
Fig. 2.16: Hydrogen Bonds can form between water and other polar molecules.
Hydrogen bond
Na Sodium atom (an uncharged
atom)
Cl Chlorine atom (an uncharged
atom)
Na+ Sodium ion (a cation)
Cl– Chlorine ion (an anion)
Sodium chloride (NaCl)
Fig. 2.14: Ionic Bonds
+ -
Solvation: Water molecules surround ions and neutralize their charge.
Sodium (Na) Chlorine (Cl)
Weak Chemical Bonds: van der Waals Interactions
δ- δ+
δ- δ+ δ- δ+ δ- δ+
δ- δ+ δ- δ+
Weak attraction of atoms due to constant motion of electrons:
High Energy Very Low Energy Low Energy
Weak attraction of atoms due to the constant motion of electrons
Chemical Bonds and Interactions
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